WEEK 1

TOPIC: INTRODUCTION TO CHEMISTRY AND ATOMIC THEORY

1. Introduction to Chemistry

Definition

Chemistry is the branch of science that studies the composition, structure, properties, and changes of matter.

Matter is anything that has mass and occupies space — air, blood, water, bones, etc.

Everyday Example:

1.      Chemistry is like cooking:

• Ingredients (chemicals) are mixed in specific ratios.
• The heat (energy) causes them to transform into something new — just like a chemical reaction!

2.      In the human body, chemistry is the language of life.

• The digestion of food, metabolism of drugs, respiration, and nerve transmission are all chemical processes.
• For instance, the conversion of glucose to energy (ATP) in cells follows the same principles as chemical reactions studied in chemistry.

 

Branches of chemistry:

General/Analytical chemistry (composition/properties/reaction)

Inorganic chemistry (compounds that do not contain Carbon)

Organic chemistry (compounds containing Carbon)

Physical chemistry (physical nature of matter)

Chemistry and other disciplines

Chemistry cuts across many disciplines of study such as:

Geology (geochemistry- soil, ores, minerals)

Biology (biochemistry- chemistry of proteins)

Agriculture (phytochemistry (antioxidants, proximate analysis)

Pharmacy

(pharmaceutical chemistry, pharmacokinetics) etc

2. Importance of Chemistry in Health and Medicine

1. Pharmacology: Understanding how drugs react in the body (drug metabolism, absorption, and excretion).
2. Biochemistry: Explains how enzymes, hormones, and DNA function chemically.
3. Pathology: Chemical tests help diagnose diseases (e.g., blood glucose, cholesterol levels).
4. Anesthesia and Radiology: Depend on chemical principles of gases and reactions.
5. Sterilization: Use of chemical disinfectants (like ethanol and chlorine).

 

What are Scientific Methods? - a set of general principles that describes how any investigation is done. The method follows a specific sequence as indicated below:

6.      Observation – describes and measures a data

7.      Hypothesis – possible interpretation of the data (must be stated in a way that it can be verified or tested (e.g. If I say that putting my wet clothes in a dryer, the clothes will be dry in 30 minutes. The statement can be tested or verified by anybody. A hypothesis is a tentative explanation of observations that acts as a guide for gathering and checking information.

8.      Experiments- test that is done to verified the hypothesis. We test a hypothesis by experimentation, calculation, and/or comparison with the experiments of others and then refine it as needed.

9.      Theory – If a hypothesis turns out to be capable of explaining a large body of experimental data, it can reach the status of a theory. Scientific theories are well-substantiated, comprehensive, testable explanations of particular aspects of nature. Theories are accepted because they provide satisfactory explanations, but they can be modified if new data become available. The path of discovery that leads from question and observation to law or hypothesis to theory consistent data from experiments that confirms the hypothesis. Then the hypothesis becomes a theory

10.  Law- When a theory remains true over a very long period of time, with unquestionable accuracy, it becomes a law. A statement that describes a natural phenomenon that always holds true under the same conditions. Unlike theories (which explain “why”), laws describe “what happens.” ‘One of such laws is ‘the law of gravity’.

 

Phases and Classification of Matter

Matter is defined as anything that occupies space and has mass, and it is all around us. Solids and liquids are more obviously matter: We can see that they take up space, and their weight tells us that they have mass. Gases are also matter; if gases did not take up space, a balloon would

Solids, liquids, and gases are the three states of matter commonly found on earth. A solid is rigid and possesses a definite shape. A liquid flows and takes the shape of a container, except that it forms a flat or slightly curved upper surface when acted upon by gravity. (In zero gravity, liquids assume a spherical shape.) Both liquid and solid samples have volumes that are very nearly independent of pressure. A gas takes both the shape and volume of its container.

The three most common states or phases of matter are solid, liquid, and gas.

 

A fourth state of matter, plasma, occurs naturally in the interiors of stars. A plasma is a gaseous state of matter that contains appreciable numbers of electrically charged particles. The presence of these charged particles imparts unique properties to plasmas that justify their classification as a state of matter distinct from gases. In addition to stars, plasmas are found in some other high-temperature environments (both natural and man-made), such as lightning strikes, certain television screens, and specialized analytical instruments used to detect trace amounts of metals.

 

 

Atoms and Molecules

An atom is the smallest particle of an element that has the properties of that element and can enter into a chemical combination. Consider the element gold, for example. Imagine cutting a gold nugget in half, then cutting one of the halves in half, and repeating this process until a piece of gold remained that was so small that it could not be cut in half (regardless of how tiny your knife may be).

This minimally sized piece of gold is an atom (from the Greek atomos, meaning “indivisible”) This atom would no longer be gold if it were divided any further.

An atom is so small that its size is difficult to imagine. One of the smallest things we can see with our unaided eye is a single thread of a spider web:

 

These strands are about 1/10,000 of a centimeter (0.00001 cm) in diameter. Although the cross-section of one strand is almost impossible to see without a microscope, it is huge on an atomic scale. A single carbon atom in the web has a diameter of about 0.000000015 centimeter, and it would take about 7000 carbon atoms to span the diameter of the strand. To put this in perspective, if a carbon atom were the size of a dime, the cross-section of one strand would be larger than a football field, which would require about 150 million carbon atom “dimes” to cover it.

The elements hydrogen, oxygen, phosphorus, and sulfur form molecules consisting of two or more atoms of the same element. The compounds water, carbon dioxide, and glucose consist of combinations of atoms of different elements

Classifying Matter

We can classify matter into several categories. Two broad categories are mixtures and pure substances.

 A pure substance has a constant composition. All specimens of a pure substance have exactly the same makeup and properties. Any sample of sucrose (table sugar) consists of 42.1% carbon, 6.5% hydrogen, and 51.4% oxygen by mass. Any sample of sucrose also has the same physical properties, such as melting point, color, and sweetness, regardless of the source from which it is isolated.

Why This Matters in Medicine & Health

1.      Drugs as Pure Substances

o    Medications contain active pharmaceutical ingredients (APIs) that must be chemically pure.

o    Example: Pure paracetamol (acetaminophen) has the same pain-relieving effect, regardless of the manufacturer.

2.      Diagnostic Tests

o    Pure reagents (like glucose solutions for blood sugar testing) are needed to ensure accurate results.

3.      Nutrition

o    Glucose and sucrose are pure substances with fixed energy values (4 kcal/g), which helps in diet planning and energy balance studies.

4.      Toxicology

o    Impurities in what should be a pure drug can cause adverse health effects (e.g., contaminated cough syrups with toxic compounds).

Medical Example

·         If a nurse dissolves pure sodium chloride (NaCl) in water, the resulting solution is safe as IV saline (0.9%).

·         But if the salt contains impurities (like heavy metals), it could be toxic to patients.

 

We can divide pure substances into two classes: elements and compounds.

Elements: are pure substances that cannot be broken down into simpler substances by chemical changes. Examples Iron, silver, gold, aluminum, sulfur, oxygen, and copper are familiar examples.

Compounds: are pure substances that can be broken down by chemical changes. This breakdown may produce either elements or other compounds, or both. Mercury(II) oxide, an orange, crystalline solid, can be broken down by heat into the elements mercury and oxygen. When heated in the absence of air, the compound sucrose is broken down into the element carbon and the compound water

(The initial stage of this process, when the sugar is turning brown, is known as caramelization—this is what imparts the characteristic sweet and nutty flavor to caramel apples, caramelized onions, and caramel). Silver(I) chloride is a white solid that can be broken down into its elements, silver and chlorine, by absorption of light. This property is the basis for the use of this compound in photographic films and photochromic eyeglasses (those with lenses that darken when exposed to light).

The properties of combined elements are different from those in the free, or uncombined, state. For example, white crystalline sugar (sucrose) is a compound resulting from the chemical combination of the element carbon, which is a black solid in one of its uncombined forms, and the two elements hydrogen and oxygen, which are colorless gases when uncombined. Free sodium, an element that is a soft, shiny, metallic solid, and free chlorine, an element that is a yellow-green gas, combine to form sodium chloride (table salt), a compound that is a white, crystalline solid.

 

Mixture

A mixture is composed of two or more types of matter that can be present in varying amounts and can be separated by physical changes, such as evaporation. There are two types of mixtures;

 A mixture with a composition that varies from point to point is called a heterogeneous mixture. Examples of heterogeneous mixtures are chocolate chip cookies, muddy water and granite.

 Medical Example: Whole blood is a heterogeneous mixture—you can separate plasma, white cells, and red cells by centrifugation.

A homogeneous mixture, also called a solution, exhibits a uniform composition and appears visually the same throughout. An example of a solution is a soft drinks, consisting of water, sugar, coloring, flavoring, and electrolytes mixed together uniformly. Each drop of a soft drink tastes the same because each drop contains the same amounts of water, sugar, and other components. Other examples of homogeneous mixtures include air, maple syrup, gasoline, and a solution of salt in water.

Medical Example:

·         IV saline solution (0.9% NaCl) is homogeneous—every drop has the same concentration of salt.

·         Blood plasma is a homogeneous mixture of water, proteins, electrolytes, and nutrients.

Why This Matters in Medicine & Health

·         IV Therapy → Patients need homogeneous solutions (saline, glucose solutions) so each drop delivers the same dose.

·         Diagnostics → Blood and urine are mixtures; lab tests often separate or analyze their components.

·         Pharmacology → Syrups, injections, and inhalers are often mixtures (solutions or suspensions). Correct preparation ensures uniformity.

·         Nutrition → Foods can be homogeneous (milk after processing) or heterogeneous (salads, cereals), which affects digestion and absorption.

 

Quick Comparison: Pure Substances vs. Mixtures

Feature	.Pure Substance	Mixture
Composition	Fixed	Variable
Properties	Constant	Depend on ratio of components
Separation	Chemical methods needed	Physical methods (e.g., filtration, distillation)
Examples	Water (H₂O), glucose, oxygen	Air, blood, saltwater, IV fluids

 

 

 

 

A summary of how to distinquish between the various major classifications of matter is shown below:

 

 

Elemental Composition of Earth

 

Element	Symbol	Percent Mass		Element	Symbol	Percent Mass
oxygen	O	49.20		chlorine	Cl	0.19
silicon	Si	25.67		phosphorus	P	0.11

 

aluminum	Al	7.50		manganese	Mn	0.09
iron	Fe	4.71		carbon	C	0.08
calcium	Ca	3.39		sulfur	S	0.06
sodium	Na	2.63		barium	Ba	0.04
potassium	K	2.40		nitrogen	N	0.03
magnesium	Mg	1.93		fluorine	F	0.03
hydrogen	H	0.87		strontium	Sr	0.02
titanium	Ti	0.58		all others	-	0.47

Physical and Chemical Properties

Physical property/Physical Change: The characteristics that enable us to distinguish one substance from another are called properties. A physical property is a characteristic of matter that is not associated with a change in its chemical composition. Familiar examples of physical properties include density, color, hardness, melting and boiling points, and electrical conductivity. We can observe some physical properties, such as density and color, without changing the physical state of the matter observed. Other physical properties, such as the melting temperature of iron or the freezing temperature of water, can only be observed as matter undergoes a physical change. A physical change is a change in the state or properties of matter without any accompanying change in its chemical composition (the identities of the substances contained in the matter). Examples:

Wax melting (still wax, just liquid);
Ice melting into water (still H2O);
Sugar dissolving in coffee (no new substance formed);
Grinding a tablet into powder for easy swallowing;
Steam condensing into liquid water.

Other examples of physical changes include magnetizing and demagnetizing metals (as is done with common antitheft security tags) and grinding solids into powders (which can sometimes yield noticeable changes in color). In each of these examples, there is a change in the physical state, form, or properties of the substance, but no change in its chemical composition.

 

 

 

Property/Change	Definition	Examples	Medical Relevance
Physical Property	Can be observed without changing substance’s identity	Color, density, boiling point	Blood color, saline conductivity
Physical Change	Change in state/form without new substance	Melting ice, dissolving sugar	Crushing tablets, nebulization
Chemical  Property	How a substance reacts chemically	Flammability, acidity	Drug reactivity, enzyme activity
Chemical Change	New substance formed	Burning wood, souring milk	Digestion, drug metabolism, sterilization

Extensive and Intensive Properties of Matter   

 

Properties of matter fall into one of two categories. If the property depends on the amount of matter present, it is an extensive property. The mass and volume of a substance are examples of extensive properties; for instance, a gallon of milk has a larger mass and volume than a cup of milk. The value of an extensive property is directly proportional to the amount of matter in question. Extensive Properties

·         Depend on the amount of matter in the sample.

·         Examples:

o    Mass → A bag of IV saline weighs more than a single IV ampoule.

o    Volume → 1 liter of blood plasma vs. 5 mL of plasma.

o    Heat content (energy) → A large pot of hot soup contains more total heat than a single spoonful at the same temperature.

·         Key Point: The value changes if the sample size changes.

Medical Relevance:

·         The volume of blood loss is an extensive property—losing 1 liter of blood is life-threatening, while losing a few milliliters is usually harmless.

·         The dose of a drug is an extensive property—the amount administered matters (e.g., 500 mg vs. 1 g of paracetamol).

If the property of a sample of matter does not depend on the amount of matter present, it is an intensive property. Temperature is an example of an intensive property. If the gallon and cup of milk are each at 20 °C (room temperature), when they are combined, the temperature remains at 20 °C. As another example, consider the distinct but related properties of heat and temperature. A drop of hot cooking oil spattered on your arm causes brief, minor discomfort, whereas a pot of hot oil yields severe burns. Both the drop and the pot of oil are at the same temperature (an intensive property), but the pot clearly contains much more heat (extensive property).

 

Medical & Health Applications

·         Temperature vs. Fever

o    Temperature (intensive): A patient’s body temperature of 39 °C is fever, whether the person weighs 40 kg or 100 kg.

o    Heat content (extensive): Larger bodies store more heat energy at the same temperature, which is why obese patients may retain heat differently.

·         Drug Concentration vs. Dose

o    Concentration in blood (intensive): 10 mg/L of a drug is the same concentration regardless of total blood volume.

o    Dose (extensive): 500 mg vs. 1000 mg paracetamol → the effect depends on the total amount given.

·         Blood Volume vs. Blood Pressure

o    Blood volume (extensive): 1 L vs. 5 L is not the same.

o    Blood pressure (intensive): 120/80 mmHg describes the state of pressure regardless of how much blood is in circulation.

SUMMARY:

·         Intensive = Quality (doesn’t change with size).

·         Extensive = Quantity (depends on how much you have).

 

6. Chemical Reactions

Definition

A chemical reaction is a process where one or more substances (reactants) are transformed into new substances (products).

Example:

2H₂ + O₂ →2H₂O

Examples:

1. Cooking an egg — once it’s cooked, it cannot turn back into raw egg.
That’s a chemical change (irreversible reaction).

2.Cellular Respiration:

C₆H₁₂ O₆ +6O₂→ 6CO₂ + 6H₂O + Energy(ATP)

— how your body converts glucose into usable energy.

1. Drug Metabolism (in the liver):
   Paracetamol is chemically broken down by enzymes into harmless metabolites.
2. Photosynthesis (in plants):
   Provides the oxygen we breathe and food we eat.
3. Acid–Base Reactions:
   Maintaining blood pH (7.35–7.45) depends on chemical buffers like bicarbonate (HCO₃⁻) reacting with acids and bases.

7. Types of Chemical Reactions

Type	Description	Example	Medical / Everyday Analogy
Combination	Two or more substances combine to form one	2H₂ + O₂ → 2H₂O	Formation of water in the body during metabolism
Decomposition	A compound breaks down into simpler substances	2H₂O → 2H₂ + O₂	Breakdown of glucose in cellular respiration
Single Displacement	One element replaces another in a compound	Zn + H₂SO₄ → ZnSO₄ + H₂	Displacement in drug–metal interactions
Double Displacement	Exchange of ions between two compounds	NaCl + AgNO₃ → NaNO₃ + AgCl	Precipitation in diagnostic tests (e.g. AgCl in chloride test)
Combustion	Reaction with oxygen releasing energy	CH₄ + 2O₂ → CO₂ + 2H₂O	Burning of glucose in cells for energy
Neutralization	Acid + Base → Salt + Water	HCl + NaOH → NaCl + H₂O	Antacids neutralizing excess stomach acid

8. Summary

• Chemistry explains how matter behaves and how changes occur.
• Atoms are building blocks of matter.
• Molecules form when atoms combine.
• Chemical reactions drive life processes — digestion, respiration, drug metabolism, and detoxification.
• Understanding atomic theory helps in radiology, pharmacology, toxicology, and biochemistry.

 

 

 

ATOMIC THEORY (Historical Background)

a) Dalton’s Atomic Theory (1808)

1. Matter is made up of indivisible particles called atoms.
2. Atoms of the same element are identical; different elements have different atoms.
3. Atoms combine in simple whole-number ratios to form compounds.
4. Chemical reactions involve rearrangement of atoms — atoms are not created or destroyed.

Limitation: It couldn’t explain isotopes or subatomic particles.

 

b) Thomson’s Model (1898 – Plum Pudding Model)

• Atom is a positively charged sphere with negatively charged electrons embedded within it.
  Analogy: Like raisins in a pudding or chocolate chips in a cookie.

 

c) Rutherford’s Nuclear Model (1911)

• Most of the atom’s mass is concentrated in a small, dense nucleus.
• Electrons move around the nucleus in mostly empty space.

Experiment: Gold foil experiment.

 

d) Bohr’s Model (1913)

• Electrons orbit the nucleus in specific energy levels (shells).
• Energy is absorbed or emitted when electrons jump between shells.

 

e) Quantum Mechanical Model

• Modern view: Electrons exist in regions called orbitals, not fixed paths.
• Predicts probabilities of finding an electron around the nucleus.

 

Modern Electronic Theory of Atoms

1. Introduction

The modern electronic theory of atoms explains how electrons are arranged and behave within atoms using quantum mechanics.

Earlier models (like Dalton’s, Thomson’s, Rutherford’s, and Bohr’s) described the atom’s structure but could not explain complex electron behavior — especially why atoms emit light in certain colors or how chemical bonds form.

Quantum mechanics came to solve these mysteries.

 

2. Development of Modern Atomic Theory

a) Bohr’s Model (1913)

• Proposed that electrons move around the nucleus in fixed energy levels (shells).
• Energy is absorbed or released when electrons jump between these levels.

Equation:

E=hνE = h\nuE=hν

Where:

• E = energy
• h = Planck’s constant
• ν = frequency of radiation

Limitation:
Bohr’s model only explained hydrogen (one-electron atom) — not atoms with multiple electrons.

 

b) Quantum Mechanical Model (Modern View)

Developed by Erwin Schrödinger, Werner Heisenberg, and others.
It treats electrons not as particles moving in definite orbits but as waves existing in regions of probability.

Key idea:
We cannot pinpoint where an electron is, but we can predict where it is most likely to be found — this region is called an orbital.

 

3. Important Principles of Quantum Mechanics

1. Wave-Particle Duality (de Broglie)

• Electrons behave both as particles (with mass) and waves (with wavelength).
• Just like light, which behaves as both a wave and a particle.

Medical example:
In X-rays and MRI, radiation (like electrons or photons) behaves both as particles and waves, allowing medical imaging.

2. Heisenberg’s Uncertainty Principle

It is impossible to know both the exact position and momentum of an electron at the same time.

Medical example:
In nuclear medicine, particle–wave uncertainty helps explain radiation spread during imaging — we can predict where particles might be, but not with exact certainty.

 

3. Schrödinger’s Wave Equation

• Describes the behavior of an electron as a wave.
• The solution to this equation gives orbitals, regions in space where electrons are most likely to be found.

 

4. Quantum Numbers

Quantum numbers are four numbers that describe the position and energy of an electron in an atom — like a unique address for each electron.

Quantum Number	Symbol	Describes	Possible Values	Example
Principal	n	Main energy level (shell)	1, 2, 3, 4…	n = 1 (K shell), n = 2 (L shell)
Azimuthal / Angular Momentum	l	Sublevel or shape of orbital	0 → (n–1)	l = 0 (s), 1 (p), 2 (d), 3 (f)
Magnetic	mₗ	Orientation of orbital	–l → +l	For l = 1, mₗ = –1, 0, +1
Spin	mₛ	Spin of electron	+½ or –½	Two possible spins per orbital

 

Think of a hospital system:

• n (principal number): The hospital building (energy level)
• l (sublevel): The department (ward or clinic)
• mₗ (orbital): The specific bed/room where a patient (electron) stays
• mₛ (spin): Whether the patient faces up or down (↑ or ↓)

Each electron has a unique “hospital address” inside the atom!

 

5. Shapes of Orbitals

Type	Symbol	Shape	Maximum Electrons	Analogy
s-orbital	s	Spherical	2	Like a round football field
p-orbital	p	Dumbbell	6	Like two balloons tied together
d-orbital	d	Cloverleaf	10	Complex, like flower petals
f-orbital	f	Very complex	14	Multilobed — rarely in basic atoms

 

 

 

According to the quantum atomic model, an atom can have many possible numbers of orbitals. These orbitals can be categorized on the basis of their size, shape or orientation. A smaller sized orbital means there is a greater chance of getting an electron near the nucleus. The orbital wave function or ϕ is a mathematical function used for representing the coordinates of an electron. The square of the orbital wave function represents the probability of finding an electron.

This wave function also helps us in drawing boundary surface diagrams. Boundary surface diagrams of the constant probability density for different orbitals help us understand the shape of orbitals.

Let us represent the shapes of orbitals with the help of boundary surface diagrams:

The Shape of s Orbitals

• The boundary surface diagram for the s orbital looks like a sphere having the nucleus as its centre which in two dimensions can be seen as a circle.
• Hence, we can say that s-orbitals are spherically symmetric having the probability of finding the electron at a given distance equal in all the directions.
• The size of the s orbital is also found to increase with the increase in the value of the principal quantum number (n), thus, 4s > 3s> 2s > 1s.

 

The Shape of p Orbitals

• Each p orbital consists of two sections better known as lobes which lie on either side of the plane passing through the nucleus.
• The three p orbitals differ in the way the lobes are oriented whereas they are identical in terms of size, shape, and energy.
• As the lobes lie along one of the x, y or z-axis, these three orbitals are given the designations 2p_x, 2p_y, and 2p_z. Thus, we can say that there are three p orbitals whose axes are mutually perpendicular.
• Similar to s orbitals the size, and energy of p orbitals increase with an increase in the principal quantum number (4p > 3p > 2p).

The Shape of p Orbitals

• Each p orbital consists of two sections better known as lobes which lie on either side of the plane passing through the nucleus.
• The three p orbitals differ in the way the lobes are oriented whereas they are identical in terms of size, shape, and energy.
• As the lobes lie along one of the x, y or z-axis, these three orbitals are given the designations 2p_x, 2p_y, and 2p_z. Thus, we can say that there are three p orbitals whose axes are mutually perpendicular.
• Similar to s orbitals the size, and energy of p orbitals increase with an increase in the principal quantum number (4p > 3p > 2p).

The Shape of d Orbitals

• The magnetic orbital quantum number for d orbitals is given as (-2,-1,0, 1,2). Hence, we can say that there are five d-orbitals.
• These orbitals are designated as d_xy, d_yz, d_xz, d_x²_–y ² and d_z².
• Out of these five d orbitals, the shapes of the first four d-orbitals are similar to each other, which is different from the d_z² orbital whereas the energy of all five d orbitals is the same.

6. Electronic Configuration

Definition

The arrangement of electrons in various orbitals of an atom.

Rules: Electrons are arranged in sub-orbital by three principles, viz:

Aufbau or building principles: electrons are filled into orbitals in order of increasing energy- 1s<2s<2p<3s<sp<4s<3d<4p<5s<4d<5p<6s<4f<5d etc

Pauli’s exclusion principle: Electrons spins on its axis and hence generates magnetic and electric fields. Electrons in the same degenerate orbital repels each other and therefore will have opposite spins number to be able to reside in the same orbital. Hence Pauli formulate the following rules: (i)The maximum number of electrons in a degenerate orbital is two (ii) the two electrons in a degenerate orbital have opposite quantum spin number (+/-). Each spin is represented by half headed arrows (  ). (iii) no two electron have all four principal quantum numbers the same

Hund’s rule (the principle of maximum multiplicity): degenerate orbitals are singly filled before pairing up. As a result, an atom tends to have as many unpaired electrons as possible, In order words, electrons avoid repelling each other by seeking out empty shells to fill up instead of pairing up.

​The electronic arrangement can also be written in the format shown below by following the direction of the arrow

 

 

 

1s 2s

2p

3s

3p

3d

4s

4p

4d

4f

5s

5p

5d

5f

6s

6p

6d

6f

​

​

 

​

 

​

 

​

 

​

 

Electronic Configuration and Electronic Diagram

The shorthand form of representing the arrangement of electrons in orbitals can be in form of a diagram or in a specified configuration.

For example, carbon has 6 valence electrons and can be represented thus:

Electronic Configuration: 1s2 2s22p2

The correct arrangement of the sub-orbitals is show in the previous slide

Electronic Diagram: The electronic diagram of carbon can be drawn as shown below

​

 

2s

2p

1s

 

Abbreviated Format

Another way of writing electronic configuration is to use the Abbreviated format. This involves using the configuration of noble gas preceding the desired element and filling the remaining valence electrons in the next principal quantum level

For example, Sodium has 11 valence electrons: 1s2 2s22p63s1. The nearest noble gas before sodium is Neon with electronic configuration: 1s2 2s22p61s2 2s22p6 which is one electron shorn of sodium configuration. Hence sodium configuration can be written as: [Ne]3s1 which is the condensed format of sodium configuration.

Examples

1. Hydrogen (Z = 1): 1s¹

 

2. Helium (Z = 2): 1s²

 

 

3. Carbon (Z = 6): 1s² 2s² 2p²

 

4. Sodium (Z = 11): 1s² 2s² 2p⁶ 3s¹

 

5. Calcium (Z = 20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

 

7. Periodic Properties and Atomic Structure

The arrangement of electrons explains periodic trends:

Property	Definition	Trend	Medical Analogy
Atomic Radius	Size of an atom	Decreases across a period, increases down a group	Small drug molecules penetrate cells more easily
Ionization Energy	Energy to remove an electron	Increases across a period	Similar to how tightly bound oxygen is to hemoglobin
Electronegativity	Tendency to attract electrons	Increases across a period	Like the “attraction” between receptor and ligand in drug binding
Electron Affinity	Energy change when atom gains an electron	Increases across a period	Like enzymes capturing substrates more efficiently

 

 

 

 

 

Medical Relevance of Modern Atomic Theory

1. Drug–Receptor Interaction:
   Electron configuration determines how drugs bind to receptors — e.g., hydrogen bonding and van der Waals forces depend on electron density.
2. Radiation and Imaging:
• X-rays and CT scans depend on the interaction of high-energy electrons with tissues.
• Understanding atomic structure helps explain ionizing radiation effects on cells.
3. Isotopes in Medicine:
• Cobalt-60 and Iodine-131 used in radiotherapy and diagnosis are based on the concept of atomic nuclei and electron transitions.
4. Enzyme Catalysis:
• The reaction rates in biochemistry depend on how electrons are shared or transferred between molecules (redox reactions).

 

 Summary

• The modern atomic theory uses quantum mechanics to describe how electrons behave as waves and particles.
• Orbitals represent regions where electrons are likely to be found.
• Quantum numbers act like an address for each electron.
• Electron configuration determines the chemical behavior of elements.
• These concepts explain biochemical reactions, drug action, and radiation therapy in medicine.

 

10. Review Questions

1. What are the key differences between Bohr’s and Schrödinger’s atomic models?
2. Explain wave-particle duality and its medical relevance.
3. List and describe the four quantum numbers.
4. Write the electron configurations of Na, Cl, and Fe.
5. State and explain the three rules governing electron filling.
6. How does atomic structure explain the action of radioactive isotopes in medicine.

Electronic Configuration and the Periodic Table

 

1. Introduction

The Periodic Table is the chemist’s map of elements — it organizes all known elements according to their atomic structure and electronic configuration.

The arrangement helps us:

• Predict how elements behave in chemical reactions.
• Understand their valence electrons, bonding capacity, and biological roles.

In medicine and biochemistry, this knowledge helps explain:

• How ions like Na⁺, K⁺, and Ca²⁺ function in the body.
• How transition metals like Fe, Cu, and Zn act as cofactors in enzymes.
• How radioactive elements behave in diagnostic imaging and radiation therapy.

·         Classification of Elements

·         Many elements differ dramatically in their chemical and physical properties, but some elements are similar in their behaviors. We can sort the elements into large classes with common properties:

·         Metals: Elements that are shiny, malleable, ductile,  good conductors of heat and electricity.

·         Nonmetals: Elements that appear dull, poor conductors of heat and electricity;

·         Metalloids: Elements that conduct heat and electricity moderately well, and possess some properties of metals and some properties of nonmetals.

·          

·         The elements can also be classified based on their position in the periodic table

·          The main-group elements (or representative elements) in the columns labeled 1, 2, and 13–18; The transition metals in the columns labeled 3–12; and

·          Inner transition metals in the two rows at the bottom of the table (the top-row elements are called lanthanides and the bottom-row elements are actinides.

·          

·         The elements can be subdivided further by more specific properties, such as the composition of the compounds they form:

·         Alkali metals:The elements in group 1 (the first column) form compounds that consist of one atom of the element and one atom of hydrogen. and they all have similar chemical properties.

·          Alkaline earth metals:  The elements in group 2 (the second column) form compounds consisting of one atom of the element and two atoms of hydrogen, with similar properties among members of that group.

·          The pnictogens (group 15)

·         The Chalcogens (group 16)

·         The Halogens (group 17),

·         The Noble gases (group 18 or zero, also known as inert or rare gases).

Electronic Configuration and the Periodic Table

a) Periodic Law

The properties of elements are periodic functions of their atomic numbers — that is, elements show repeating patterns when arranged by increasing atomic number.

b) Structure of the Periodic Table

The modern periodic table is arranged into:

• Periods (rows): 7 horizontal rows representing energy levels (n = 1–7)
• Groups (columns): 18 vertical columns showing elements with similar valence electron configurations

c) Blocks of the Periodic Table

The table is divided based on the type of orbital being filled:

Block	Sublevel Being Filled	Example Elements	Range of Groups	Example Configuration
s-block	s orbital	H, He, Li, Na, Mg	1–2	ns¹–²
p-block	p orbital	C, N, O, F, Cl	13–18	ns² np¹–⁶
d-block	d orbital (Transition metals)	Fe, Cu, Zn	3–12	(n–1)d¹–¹⁰ ns⁰–²
f-block	f orbital (Lanthanides/Actinides)	Ce, U	Inner transition	(n–2)f¹–¹⁴ (n–1)d⁰–¹   ns²

 

a) Groups (Columns)

Elements in the same group have:

• Same number of valence electrons
• Similar chemical and physical properties

Group	Name	Valence Electrons	Example	Medical Relevance
1	Alkali metals	1	Na, K	Important for nerve impulses, fluid balance
2	Alkaline earth metals	2	Ca, Mg	Bone formation, muscle contraction
17	Halogens	7	Cl, I	Used as disinfectants and in thyroid hormones
18	Noble gases	8	He, Ne, Ar	Inert gases, helium used in cryogenics and MRI cooling

 

b) Periods (Rows)

• Each period represents a new energy level (shell) being filled.
• Properties change gradually across a period due to change in effective nuclear charge and valence electron configuration.

 

5. Periodic Trends and Their Significance

Property	Trend Across a Period	Trend Down a Group	Medical/Everyday Examples
Atomic Radius	Decreases (nucleus pulls electrons closer)	Increases (more shells added)	Like tightening or loosening a belt — atoms “shrink” or “expand”
Ionization Energy (Energy to remove one electron)	Increases	Decreases	Like removing oxygen from hemoglobin — tightly bound in some, loosely in others
Electronegativity (Attraction for electrons)	Increases	Decreases	Like magnet strength — stronger across a period
Electron Affinity	Increases	Decreases	Similar to how receptors “attract” certain ligands more strongly
Metallic Character	Decreases	Increases	Explains why Na (metal) is reactive but Cl (nonmetal) forms salts

 

Example of Periodic Behavior

• Sodium (Na): 1s²2s²2p⁶3s¹ → reactive metal, donates electrons easily.
• Chlorine (Cl): 1s²2s²2p⁶3s²3p⁵ → reactive nonmetal, gains electrons easily.
• Together, they form NaCl (table salt), essential for maintaining electrolyte balance in the body.

6. Special Groups and Their Biological Importance

Element/Group	Function in the Body
Na⁺ & K⁺ (Group 1)	Maintain osmotic balance and nerve transmission
Ca²⁺ & Mg²⁺ (Group 2)	Bone and muscle function, enzyme activation
Fe, Cu, Zn (Transition Metals)	Components of hemoglobin, cytochromes, and enzymes
I (Halogen)	Part of thyroid hormones (T₃, T₄) regulating metabolism
C, H, O, N (p-block)	Basic building blocks of biomolecules (proteins, carbs, lipids)

 

7. Relationship Between Electronic Configuration and Chemical Behavior

The valence electrons (outermost electrons) determine how atoms bond or react.

Example	Configuration	Behavior
Na (Z=11): 1s²2s²2p⁶3s¹	Loses 1 electron → Na⁺	Metal (donor)
Cl (Z=17): 1s²2s²2p⁶3s²3p⁵	Gains 1 electron → Cl⁻	Nonmetal (acceptor)
O (Z=8): 1s²2s²2p⁴	Gains 2 electrons → O²⁻	Forms oxides, vital for energy release
Ca (Z=20): [Ar]4s²	Loses 2 electrons → Ca²⁺	Essential for muscle contraction

 

 

8. Everyday and Medical Relevance

1. Electrolyte Balance:
   Na⁺, K⁺, Ca²⁺, and Cl⁻ ions maintain nerve impulses, heart rhythm, and fluid balance.
2. Oxygen Transport:
   Fe²⁺/Fe³⁺ ions in hemoglobin depend on transition metal chemistry.
3. Drug Design:
   Electronic structure predicts how drugs interact with enzymes and receptors.
4. Radiotherapy:
   Elements like Cobalt-60 (transition metal isotope) used for cancer treatment are chosen for their electronic and nuclear properties.
5. Biochemical Catalysis:
   Zn²⁺ in carbonic anhydrase and Mg²⁺ in ATP reactions rely on electron configuration for proper catalytic function.

 

9. Step-by-Step Summary on the Periodic Table

 

1.    Arrangement by Atomic Mass:

 In the early stages of the periodic table's development, elements were arranged in order of increasing atomic mass.  Chemists observed that certain properties of elements repeated periodically as their atomic masses increased.

2.    Observation of Periodic Trends:

 As elements were arranged by atomic mass, scientists started to notice patterns in their properties.

 Certain properties, such as atomic radius, electronegativity, and ionization energy, showed periodic trends, meaning they varied in a regular and predictable manner across the elements.

3.    Mendeleev's Periodic Law:

 Dmitri Mendeleev, a Russian chemist, formulated the Periodic Law in 1869.

 Mendeleev arranged the elements in order of increasing atomic mass and grouped them based on their similar chemical properties.

 He left gaps for undiscovered elements and predicted their properties based on the periodic trends and the patterns he observed.

4.    Discovery of Atomic Number:

 In the early 20th century, Henry Moseley's experiments with X-ray spectroscopy revealed a fundamental property of elements called atomic number.

 Atomic number represents the number of protons in an atom's nucleus and determines an element's identity.

 Moseley's discovery led to the understanding that the periodic table should be arranged based on atomic number rather than atomic mass.

5.    Modern Periodic Table:

 The modern periodic table organizes elements based on their atomic number, with elements arranged in order of increasing atomic number from left to right and top to bottom.

 Elements are grouped into periods (horizontal rows) and groups/families (vertical columns) based on their electron configurations and similar properties.

 The periodic table consists of multiple blocks, including the s-block, p-block, d-block, and f-block, representing the different types of orbitals that electrons occupy.

6.    Further Refinements and Expansions:

 Over time, the periodic table has been refined and expanded as new elements have been discovered and synthesized.

      Elements beyond atomic number 118 are yet to be officially named and have temporary systematic names based on their atomic numbers.

The Periodic Table is divided into Several Blocks

1.      S-block: Elements in the first two groups (Group 1 and Group 2) are called s-block elements. They have their outermost electron(s) in the s orbital.

2.      P-block: Elements in groups 13 to 18, excluding helium, are known as p-block elements. They have their outermost electron(s) in the p orbital.

3.      D-block: Elements in groups 3 to 12 are called d-block elements or transition metals. They have their outermost electron(s) in the d orbital.

4.      F-block: The elements at the bottom of the periodic table, also known as inner transition metals, are called f-block elements. They have their outermost electron(s) in the f orbital. The f-block is further divided into the lanthanide and actinide series.

The periodic table provides information about an element's atomic mass, atomic radius, electronegativity, and various other properties. It also helps predict an element's reactivity, bonding patterns, and potential chemical reactions.

 

 

10. Review Questions

1. Define electronic configuration and state the three rules guiding electron filling.
2. Differentiate between periods and groups in the periodic table.
3. Write the electronic configurations for Na, Ca, and Cl, and explain their chemical behavior.
4. Explain the relationship between valence electrons and periodic trends.
5. How is the periodic table divided into s, p, d, and f blocks?
6. Mention three ways electronic configuration is important in medicine.
7. Why do noble gases show little or no reactivity?
8. Explain the biological significance of Ca²⁺ and Fe²⁺ ions in the bod

: Measurements

What is Measurement? Measurement is the assigning of numerals to objects, events, etc. Measurements provide the macroscopic information that is the basis of most of the hypotheses, theories, and laws that describe the behavior of matter and energy in both the macroscopic and microscopic domains of chemistry. Every measurement provides three kinds of information: the size or magnitude of the measurement (a number); a standard of comparison for the measurement (a unit); and an indication of the uncertainty of the measurement. While the number and unit are explicitly represented when a quantity is written, the uncertainty is an aspect of the measurement result that is more implicitly represented. 

The number in the measurement can be represented in different ways, including decimal form and scientific notation. (Scientific notation is also known as exponential notation;) For example, the maximum takeoff weight of a Boeing 777-200ER airliner is 298,000 kilograms, which can also be written as 2.98 × 105 kg. The mass of the average mosquito is about 0.0000025 kilograms, which can be written as 2.5 × 10−6 kg.

Without units, a number can be meaningless, confusing, or possibly life threatening. Suppose a doctor prescribes phenobarbital to control a patient’s seizures and states a dosage of “100” without specifying units. Not only will this be confusing to the medical professional giving the dose, but the consequences can be dire: 100 mg given three times per day can be effective as an anticonvulsant, but a single dose of 100 g is more than 10 times the lethal amount.

We usually report the results of scientific measurements in SI units. Other units can be derived from these base units. The standards for these units are fixed by international agreement, and they are called the International System of Units or SI Units (from the French, Le Système International d’Unités). SI units have been used by the United States National Institute of Standards and Technology (NIST) since 1964.

Why Units Matter in Medicine

• Without units, numbers are meaningless and dangerous.
• Case Example:
• If a doctor prescribes “100” phenobarbital without specifying units:
• 100 mg/day → therapeutic dose.
• 100 g → lethal overdose (over 1,000× too much).

Lesson: Always record number + unit.

Medical Relevance of Accurate Measurements

1. Dosage calculations – Overdose or underdose can harm patients.
2. Diagnostics – Blood glucose, cholesterol, and blood pressure depend on precise measurements.
3. IV therapy – Flow rates must be measured carefully (e.g., 50 mL/hr vs. 500 mL/hr makes the difference between safe hydration and fluid overload).
4. Research & Clinical Trials – Drug effectiveness depends on reproducible measurements.

 

Base Units of the SI System

 

Property Measured	Name of Unit	Symbol of Unit
length	meter	m
mass	kilogram	kg
time	second	s
temperature	kelvin	K
electric current	ampere	A
amount of substance	mole	mol
luminous intensity	candela	cd

 

We also use fractions or multiples of units in the SI system, but these fractions or multiples are always powers of 10. Fractional or multiple SI units are named using a prefix and the name of the base unit. For example, a length of 1000 meters is also called a kilometer because the prefix kilo means “one thousand,” which in scientific notation is 103 (1 kilometer = 1000 m = 103 m). The prefixes used and the powers to which 10 are raised are listed in the table below:

 

Common Unit Prefixes

 

Prefix	Symbol	Factor	Example
femto	f	10−15	1 femtosecond (fs) = 1 × 10−15 m (0.000000000000001 s)
pico	p	10−12	1 picometer (pm) = 1 × 10−12 m (0.000000000001 m)
nano	n	10−9	4 nanograms (ng) = 4 × 10−9 g (0.000000004 g)
micro	µ	10−6	1 microliter (μL) = 1 × 10−6 L (0.000001 L)
milli	m	10−3	2 millimoles (mmol) = 2 × 10−3 mol (0.002 mol)
centi	c	10−2	7 centimeters (cm) = 7 × 10−2 m (0.07 m)
deci	d	10−1	1 deciliter (dL) = 1 × 10−1 L (0.1 L)
kilo	k	103	1 kilometer (km) = 1 × 103 m (1000 m)
mega	M	106	3 megahertz (MHz) = 3 × 106 Hz (3,000,000 Hz)
giga	G	109	gigayears (Gyr) = 8 × 109 yr (8,000,000,000 Gyr)
tera	T	1012	5 terawatts (TW) = 5 × 1012 W (5,000,000,000,000 W)

                                     SI Base Units

 This section introduces four of the SI base units commonly used in chemistry.

Length

The standard unit of length in both the SI and original metric systems is the meter (m). A meter was originally specified as 1/10,000,000 of the distance from the North Pole to the equator. It is now defined as the distance light in a vacuum travels in 1/299,792,458 of a second. A meter is about 3 inches longer than a yard. One meter is about 39.37 inches or 1.094 yards. Longer distances are often reported in kilometers (1 km = 1000 m = 103 m), whereas shorter distances can be reported in centimeters (1 cm = 0.01 m = 10−2 m) or millimeters (1 mm = 0.001 m = 10−3 m).

Mass

The standard unit of mass in the SI system is the kilogram (kg). A kilogram was originally defined as the mass of a liter of water (a cube of water with an edge length of exactly 0.1 meter). It is now defined by a certain cylinder of platinum-iridium

alloy, which is kept in France. Any object with the same mass as this cylinder is said to have a mass of 1 kilogram. One kilogram is about 2.2 pounds. The gram (g) is exactly equal to 1/1000 of the mass of the kilogram (10−3 kg).

            This replica prototype kilogram is housed at the National Institute of Standards and Technology (NIST) in Maryland. (credit: National Institutes of Standards and Technology)

Temperature

Temperature is an intensive property. The SI unit of temperature is the kelvin (K). The IUPAC convention is to use kelvin (all lowercase) for the word, K (uppercase) for the unit symbol, and neither the word “degree” nor the degree symbol (°). The degree Celsius (°C) is also allowed in the SI system, with both the word “degree” and the degree symbol used for Celsius measurements. Celsius degrees are the same magnitude as those of kelvin, but the two scales place their zeros in different places. Water freezes at 273.15 K (0 °C) and boils at 373.15 K (100 °C) by definition, and normal human body temperature is approximately 310 K (37 °C).

 

Time

The SI base unit of time is the second (s). Small and large time intervals can be expressed with the appropriate prefixes; for example, 3 microseconds = 0.000003 s = 3 × 10−6 and 5 megaseconds = 5,000,000 s = 5 × 106 s. Alternatively, hours, days, and years can be used.

 

Derived SI Units

We can derive many units from the seven SI base units. For example, we can use the base unit of length to define a unit of volume, and the base units of mass and length to define a unit of density.

 

Volume

Volume is the measure of the amount of space occupied by an object. The standard SI unit of volume is defined by the base unit of length. The standard volume is a cubic meter (m3), a cube with an edge length of exactly one meter. To dispense a cubic meter of water, we could build a cubic box with edge lengths of exactly one meter. This box would hold a cubic meter of water or any other substance.

A more commonly used unit of volume is derived from the decimeter (0.1 m, or 10 cm). A cube with edge lengths of exactly one decimeter contains a volume of one cubic decimeter (dm3). A liter (L) is the more common name for the cubic decimeter. One liter is about 1.06 quarts.

A cubic centimeter (cm3) is the volume of a cube with an edge length of exactly one centimeter. The abbreviation cc (for cubic centimeter) is often used by health professionals. A cubic centimeter is also called a milliliter (mL) and is 1/1000 of a

Accuracy and Precision

Scientists typically make repeated measurements of a quantity to ensure the quality of their findings and to know both the precision and the accuracy of their results. Measurements are said to be precise if they yield very similar results when repeated in the same manner. A measurement is considered accurate if it yields a result that is very close to the true or accepted value. Precise values agree with each other; accurate values agree with a true value. These characterizations can be extended to other contexts, such as the results of an archery competition.

 

As mentioned earlier, the SI unit of temperature is the kelvin (K). Unlike the Celsius and Fahrenheit scales, the kelvin scale is an absolute temperature scale in which 0 (zero) K corresponds to the lowest temperature that can theoretically be achieved. The early 19th-century discovery of the relationship between a gas's volume and temperature suggested that the volume of a gas would be zero at −273.15 °C. In 1848, British physicist William Thompson, who later adopted the title of Lord Kelvin, proposed an absolute temperature scale based on this concept (further treatment of this topic is provided in this text’s chapter on gases).

The freezing temperature of water on this scale is 273.15 K and its boiling temperature 373.15 K. Notice the numerical difference in these two reference temperatures is 100, the same as for the Celsius scale, and so the linear

°C

 relation between these two temperature scales will exhibit a slope of 1 K . Following the same approach, the equations for converting between the kelvin and Celsius temperature scales are derived to be:

TK = T°C + 273.15

T°C = TK − 273.15

The 273.15 in these equations has been determined experimentally, so it is not exact. The diagram below shows the relationship among the three temperature scales. Recall that we do not use the degree sign with temperatures on the kelvin scale.