1. Scope and core ideas

General chemistry explains how matter is built, how it reacts, and how energy changes during those reactions. It is the language behind biology, physiology, pharmacology, and laboratory science.

The central idea is that macroscopic behavior comes from microscopic structure:

• Atomic structure influences bonding.

• Bonding influences molecular shape and polarity.

• Shape and polarity influence solubility, reactivity, and biological function.

• Concentration and pH influence chemical equilibria in cells and fluids.

Most introductory problems reduce to one of these tasks:

• Count particles or moles.

• Convert between mass, moles, and volume.

• Predict bonding or molecular geometry.

• Balance a reaction and apply stoichiometry.

• Use equilibrium, pH, or redox relationships to find an unknown.

2. Matter, atoms, and the periodic table

Matter and composition

Matter is anything with mass and volume. In chemistry, it is usually classified as:

• Element: one type of atom

• Compound: chemically combined elements in a fixed ratio

• Mixture: physical combination of substances

Mixtures may be:

• Homogeneous: uniform throughout

• Heterogeneous: nonuniform

Atomic structure

Atoms contain:

• Protons, with charge $+1$

• Neutrons, with charge $0$

• Electrons, with charge $-1$

Atomic number $Z$ identifies the element and equals the number of protons.

Mass number $A$ is:

where $N$ is the number of neutrons.

An ion is an atom or group of atoms with a net charge:

• Cation: positively charged, usually by losing electrons

• Anion: negatively charged, usually by gaining electrons

Isotopes

Isotopes have the same $Z$ but different neutron counts. Their chemical behavior is nearly the same because chemistry depends mainly on electrons, not neutrons.

Periodic table structure

The periodic table is arranged by increasing atomic number. Elements in the same column often have similar chemistry because they have similar valence electron configurations.

Key regions:

• Alkali metals: Group 1, very reactive metals

• Alkaline earth metals: Group 2

• Halogens: Group 17, very reactive nonmetals

• Noble gases: Group 18, chemically stable

3. Electron structure and periodic trends

Electron configuration

Electrons occupy quantized energy levels and orbitals. The main orbital types are:

• $s$ holds 2 electrons

• $p$ holds 6 electrons

• $d$ holds 10 electrons

• $f$ holds 14 electrons

The filling order is governed by energy ordering rather than simple shell number alone.

Valence electrons

Valence electrons are the outer-shell electrons most responsible for bonding and reactivity. For main-group elements, the number of valence electrons usually matches the group number pattern.

Common periodic trends

Across a period from left to right:

• Atomic radius decreases

• Ionization energy generally increases

• Electronegativity generally increases

• Metallic character decreases

Down a group:

• Atomic radius increases

• Ionization energy decreases

• Electronegativity generally decreases

• Metallic character increases

Why trends matter

These trends explain why:

• Alkali metals lose electrons easily.

• Halogens attract electrons strongly.

• Small, highly charged ions interact strongly with water.

• Highly electronegative atoms create polar bonds.

4. Bonding and molecular shape

Ionic, covalent, and metallic bonding

• Ionic bonding: electron transfer and electrostatic attraction, typically between metals and nonmetals

• Covalent bonding: electron sharing between nonmetals

• Metallic bonding: delocalized electrons among metal atoms

In biological systems, most important molecules are covalent, but ions still matter because they control charge balance, osmotic pressure, and electrical activity.

Lewis structures

Lewis structures show valence electrons as dots and bonds as shared pairs.

General procedure:

1. Count total valence electrons.

2. Choose the central atom, usually the least electronegative except hydrogen.

3. Connect atoms with single bonds.

4. Complete octets on outer atoms.

5. Put remaining electrons on the central atom.

6. Form multiple bonds if needed.

Formal charge

Formal charge helps evaluate reasonable Lewis structures:

Structures with smaller formal charges and minimal charge separation are usually preferred.

VSEPR geometry

Electron groups around a central atom repel each other and arrange to minimize repulsion.

Common shapes:

Polarity

Bond polarity depends on electronegativity difference. Molecular polarity depends on both bond polarity and geometry.

• Symmetric molecules may be nonpolar even if their bonds are polar.

• Asymmetric molecules are often polar.

Polarity strongly affects:

• Solubility

• Boiling point

• Membrane permeability

• Protein-ligand interactions

5. Naming, formulas, and reactions

Chemical nomenclature

You should be able to identify and name:

• Binary ionic compounds

• Covalent compounds with prefixes

• Acids

• Polyatomic ions

Examples:

• NaCl: sodium chloride

• CO$_2$: carbon dioxide

• HCl(aq): hydrochloric acid

• CaCO$_3$: calcium carbonate

Balancing equations

Chemical equations must satisfy conservation of atoms and charge.

Example:

Coefficients change amounts, not formulas.

Common reaction types

• Synthesis: two or more reactants form one product

• Decomposition: one compound breaks apart

• Single replacement: one element replaces another

• Double replacement: ions exchange partners

• Combustion: reaction with oxygen, often producing CO$_2$ and H$_2$O

• Acid-base: proton transfer

• Precipitation: insoluble solid forms

• Redox: electron transfer

6. Stoichiometry and limiting reagents

Stoichiometry converts between substances using balanced equation coefficients.

Mole concept

One mole contains Avogadro's number of entities:

Moles connect particle count, mass, and volume.

Conversion ladder

Typical workflow:

Limiting reagent

The limiting reagent is the reactant consumed first and therefore limits product formation.

Method:

1. Convert each reactant to moles.

2. Divide by its stoichiometric coefficient if needed.

3. The smallest available reaction extent is limiting.

Percent yield

Theoretical yield is the amount predicted by stoichiometry.

Actual yield is what is obtained experimentally.

Common stoichiometry mistakes

• Forgetting to balance the equation first

• Using mass ratios instead of mole ratios

• Confusing limiting reagent with excess reagent

• Dropping units during conversions

7. States of matter and gases

Phases and phase changes

Matter exists as solid, liquid, gas, and plasma. General chemistry focuses on the first three.

Phase changes include:

• Melting and freezing

• Vaporization and condensation

• Sublimation and deposition

Gas laws

The ideal gas law is:

where:

• $P$ is pressure

• $V$ is volume

• $n$ is moles

• $R$ is the gas constant

• $T$ is temperature in kelvin

Useful forms:

for a fixed amount of gas.

Partial pressures

For gas mixtures:

and

where $X_i$ is mole fraction.

Why gases matter in medicine

Gas laws explain:

• Oxygen delivery and ventilation

• Anesthetic partial pressures

• Carbon dioxide handling

• Barometric effects at altitude or in diving

8. Solutions and concentration

Dissolution

A solution is a homogeneous mixture of solute and solvent.

Important idea: "like dissolves like."

• Polar and ionic solutes tend to dissolve in polar solvents.

• Nonpolar solutes tend to dissolve in nonpolar solvents.

Concentration units

Common concentration measures:

• Molarity:

• Molality:

• Mass percent

• Mole fraction

• Parts per million

Dilution

For solutions where solute amount is conserved:

Colligative properties

Colligative properties depend on the number of dissolved particles, not identity.

Examples:

• Vapor pressure lowering

• Boiling point elevation

• Freezing point depression

• Osmotic pressure

Osmotic pressure is especially important in physiology because it affects water movement across membranes.

9. Thermochemistry

Thermochemistry studies energy transfer as heat in chemical and physical changes.

System and surroundings

• Exothermic: system releases heat, $q < 0$

• Endothermic: system absorbs heat, $q > 0$

First law idea

Energy is conserved. In chemistry, heat and work are common forms of energy transfer.

Heat capacity

Heat required to change temperature:

where:

• $m$ is mass

• $c$ is specific heat

• $\Delta T$ is temperature change

Enthalpy

At constant pressure, heat flow equals enthalpy change:

Hess's law

If a reaction can be written as the sum of multiple steps, the total enthalpy change is the sum of the step enthalpies.

This is useful when a direct reaction enthalpy is not given.

Bond enthalpy approximation

Rough estimate:

10. Chemical equilibrium

Many reactions are reversible and reach dynamic equilibrium.

Equilibrium constant

For a reaction:

the equilibrium expression is:

Pure solids and liquids are omitted from $K$.

Reaction quotient

The reaction quotient $Q$ has the same form as $K$ but uses current concentrations or pressures.

• If $Q < K$, the reaction shifts toward products.

• If $Q > K$, the reaction shifts toward reactants.

• If $Q = K$, the system is at equilibrium.

Le Chatelier's principle

If a system at equilibrium is disturbed, it shifts to oppose the disturbance.

Changes can involve:

• Concentration

• Pressure or volume

• Temperature

Common application

Equilibrium problems often combine:

• Stoichiometry

• ICE tables

• Approximation checks

where ICE stands for initial, change, equilibrium.

11. Acids, bases, and buffers

Acid-base chemistry is one of the most important parts of general chemistry for biology and medicine.

Definitions

Common definitions:

• Arrhenius acid: produces $H^+$ in water

• Arrhenius base: produces $OH^-$ in water

• Bronsted-Lowry acid: proton donor

• Bronsted-Lowry base: proton acceptor

Strong and weak acids/bases

• Strong acids and strong bases dissociate essentially completely.

• Weak acids and bases only partially ionize and therefore establish equilibria.

pH and pOH

At 25 C:

and

Acid dissociation constant

For a weak acid $HA$:

Weak base analog:

Buffers

A buffer resists pH change and usually contains a weak acid and its conjugate base, or a weak base and its conjugate acid.

Henderson-Hasselbalch equation:

Why buffers matter biologically

Buffers stabilize:

• Blood pH

• Intracellular pH

• Enzyme activity

• Protein charge state

Small pH shifts can change protein structure, binding, and catalysis.

pKa intuition

• If pH < pKa, the protonated form dominates.

• If pH > pKa, the deprotonated form dominates.

This is a key idea for amino acids, drug ionization, and membrane transport.

12. Kinetics and reaction rate

Thermodynamics tells whether a reaction is favorable; kinetics tells how fast it happens.

Rate

Reaction rate is the change in concentration per unit time.

Factors that affect rate:

• Concentration

• Temperature

• Surface area

• Catalysts

• Molecular orientation and activation energy

Collision theory

A reaction occurs when particles collide with:

• Enough energy to overcome activation energy

• Proper orientation

Rate laws

A typical rate law is:

where $m$ and $n$ are experimentally determined, not taken from the balanced equation unless the step is elementary.

Activation energy and catalysts

Catalysts lower activation energy and speed up reactions without being consumed.

They do not change:

• Overall $\Delta G$ or $\Delta H$

• The equilibrium constant

They only help the system reach equilibrium faster.

13. Redox and electrochemistry

Oxidation and reduction

Redox reactions involve electron transfer.

• Oxidation: loss of electrons

• Reduction: gain of electrons

Mnemonic: OIL RIG

Oxidation numbers

Oxidation numbers help track electron transfer.

Rules are based on common reference states, such as:

• Free elements have oxidation number 0

• Monatomic ions have oxidation number equal to charge

• Oxygen is usually $-2$

• Hydrogen is usually $+1$ except in metal hydrides

Balancing redox reactions

Two common methods:

• Oxidation number method

• Half-reaction method

The half-reaction method is especially useful in acidic or basic solution.

Electrochemical cells

An electrochemical cell converts chemical energy into electrical energy or vice versa.

Key quantities:

• Anode: oxidation occurs

• Cathode: reduction occurs

For a galvanic cell:

• Electrons flow from anode to cathode

• The cell can perform work spontaneously

Biological relevance

Redox chemistry underlies:

• Cellular respiration

• Photosynthesis

• Electron transport chains

• Antioxidant chemistry

• Metal ion chemistry in enzymes

14. General chemistry for biology and medicine

General chemistry becomes more useful when connected to real systems.

Water as the dominant solvent

Water is polar, hydrogen-bonding, and unusually good at stabilizing ions. That makes it the main solvent in living systems.

Consequences:

• Ions dissolve readily.

• Polar solutes are favored.

• Hydrophobic molecules cluster together.

• Hydrogen bonding affects folding and recognition.

pH and biomolecules

The charge on amino acids, peptides, and drugs depends on pH relative to pKa.

Practical effects:

• Protein solubility changes with pH.

• Enzyme activity depends on ionizable residues.

• Drug absorption depends on ionization state.

Buffers in physiology

Buffers keep body fluids within narrow pH ranges.

The bicarbonate system is the classic example:

This system is linked to respiration and renal control.

Osmosis and tonicity

Water moves across semipermeable membranes toward higher effective solute concentration.

Clinical relevance:

• Hypotonic, isotonic, and hypertonic solutions

• IV fluid selection

• Cell swelling or shrinkage

Electrolytes

Electrolytes are ions in solution, such as Na$^+$, K$^+$, Ca$^{2+}$, and Cl$^-$. They matter for:

• Nerve conduction

• Muscle contraction

• Fluid balance

• Acid-base control

Common laboratory ideas

• Concentration and dilution

• Serial dilution

• Calibration curves

• Spectrophotometry

• Standard solutions

These topics often appear in lab courses and clinical measurements.

15. Problem-solving workflow

General chemistry problems are easier when you follow a fixed sequence.

Recommended workflow

1. Identify the topic: stoichiometry, equilibrium, pH, gas law, or redox.

2. Write the balanced equation if a reaction is involved.

3. List known quantities with units.

4. Convert to moles when chemistry requires it.

5. Choose the governing relationship.

6. Solve symbolically before plugging in numbers.

7. Check units and significant figures.

8. Sanity-check the magnitude and chemical sense of the answer.

Decision cues

• If the problem asks "how much forms?", think stoichiometry.

• If it asks "what is the pH?", think acid-base equilibrium.

• If it asks "which direction does it shift?", think $Q$ vs. $K$.

• If it asks "how fast?", think kinetics.

• If it asks "is it spontaneous?", think redox or thermodynamics.

Common pitfalls

• Mixing up mass and moles

• Forgetting temperature must be in kelvin for gas laws

• Using strong-acid formulas on weak-acid problems

• Treating rate law orders as coefficients

• Ignoring phase labels in equilibrium expressions

• Losing track of charge balance in ionic equations

16. Formula summary

Core formulas

High-yield facts

• Bonding and shape control polarity.

• Polarity controls solubility and intermolecular interactions.

• Moles are the bridge between particles and measurable amounts.

• Buffers resist pH change but do not make pH immutable.

• Catalysts speed reactions but do not change equilibrium position.

• Chemical intuition improves when equations are balanced and units are tracked carefully.

Sources

• OpenStax University Physics

• Physics LibreTexts

• Halliday, Resnick, and Walker, Fundamentals of Physics

• Serway and Jewett, Physics for Scientists and Engineers

• Griffiths, Introduction to Electrodynamics

• Griffiths, Introduction to Quantum Mechanics

• Taylor, Classical Mechanics

• Parell GitHub repository