CHEM 1120 Review: Atoms, Bonding, Intermolecular Forces, Equilibrium, and Acids/Bases

Atoms and the Periodic Table

Structure of the Atom

Atoms are the fundamental units of matter, consisting of a nucleus (protons and neutrons) surrounded by electrons.

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle occupying the space around the nucleus.

• For a neutral atom, the number of protons equals the number of electrons.

The Periodic Table

The periodic table organizes elements by increasing atomic number and similar chemical properties.

• Main groups (1A-8A) indicate the number of valence electrons for main group elements.

• Transition metals are found in the center block.

• Elements are classified as metals, metalloids, or nonmetals.

Valence Electrons

Definition and Importance

Valence electrons are electrons in the outermost shell of an atom. They determine chemical bonding and reactivity.

• For main group elements, the number of valence electrons equals the group number.

• Valence electrons are involved in bonding and interactions.

Electron-Dot Symbols

Electron-dot (Lewis) symbols represent valence electrons as dots around the element symbol.

Group	1A	2A	3A	4A	5A	6A	7A	8A (Noble Gases)
Example	H·	Be··	B···	C····	N·····	O······	F·······	Ne········

Stability of Full Valence Shells

• Atoms are most stable with a full valence shell (noble gas configuration).

• Two ways to achieve a full shell: gain/lose electrons (ionic bonding) or share electrons (covalent bonding).

Ionic Compounds (Ch 3)

Formation and Properties

Ionic compounds form from the attraction between positively and negatively charged ions, typically between metals and nonmetals.

• No shared electrons; electrons are transferred.

• Not represented by Lewis structures.

• When dissolved in water, ions separate (dissociation).

Covalent Molecules (Ch 4)

Formation and Representation

Covalent bonds form when two nonmetals share electrons. These are represented by Lewis structures.

• Covalent molecules do not separate into ions when dissolved.

Bonding Patterns

Atoms form a predictable number of bonds based on their valence electrons:

Element	Common Bonds
Carbon (C)	4
Nitrogen (N)	3
Oxygen (O)	2
Halogen (F, Cl, etc.)	1
Hydrogen (H)	1

Molecular Geometry

The shape of molecules depends on the number of bonds and lone pairs:

Bonds	Lone Pairs	Geometry	Example
2	0	Linear	CO2
3	0	Trigonal planar	BF3
2	1	Bent	SO2
4	0	Tetrahedral	CH4
3	1	Pyramidal	NH3
2	2	Bent	H2O

Electronegativity and Bond Polarity

• Electronegativity is the tendency of an atom to attract electrons in a bond.

• If the difference in electronegativity is:

  • 0 – 0.4: Nonpolar covalent

  • 0.5 – 1.9: Polar covalent

  • 2.0 and above: Ionic

Molecular Polarity

• Molecules can be polar or nonpolar depending on both bond polarity and molecular shape.

• Examples: Water (H2O) is polar; carbon dioxide (CO2) is nonpolar.

Intermolecular Forces (Ch 8)

Types of Intermolecular Forces

• Dipole-dipole: Attraction between polar molecules (δ+ and δ- regions).

• London-dispersion: Present in all molecules; strength increases with molecular size.

• Hydrogen bonding: Strong dipole-dipole interaction; requires H bonded to O, N, or F and a lone pair on O, N, or F.

Effect on Physical Properties

• Stronger intermolecular forces lead to higher melting and boiling points.

Force	Strength	Characteristics
Dipole-dipole	Weak (1 kcal/mol)	Polar molecules
London dispersion	Weak (0.5–2.5 kcal/mol)	All molecules; increases with size
Hydrogen bond	Moderate (2–10 kcal/mol)	H with O, N, or F

Solubility (Ch 9)

Principles of Solubility

• Like dissolves like: Molecules with similar intermolecular forces mix well.

• Polar mixes with polar; nonpolar with nonpolar; H-bonding with H-bonding; ionic compounds dissolve in polar solvents.

Chemical Equilibrium (Ch 7)

Dynamic Equilibrium

• Some reactions are reversible and reach a state where the forward and reverse reaction rates are equal.

• At equilibrium, concentrations of reactants and products remain constant (not necessarily equal).

Equilibrium Constant ($K_{eq}$)

• Describes the ratio of product to reactant concentrations at equilibrium.

$aA + bB \rightleftharpoons mM + nN$

$K = \frac{[M]^m[N]^n}{[A]^a[B]^b}$

Le Châtelier’s Principle

• If a system at equilibrium is disturbed, it will shift to restore equilibrium.

• Changes in concentration, temperature, or pressure can shift the equilibrium position.

Change	Effect
Increase reactant	Shifts toward products
Increase product	Shifts toward reactants
Increase temperature (endothermic)	Shifts toward products
Increase pressure (gases)	Shifts toward fewer moles of gas

Energy Diagrams

• Show the energy changes during a reaction.

• A catalyst lowers the activation energy ($E_{act}$), making the reaction faster.

Acids and Bases (Ch 10)

Definitions

• Acid: Substance that donates an H+ ion.

• Base: Substance that accepts an H+ ion.

• Conjugate acid-base pairs: Differ by one H+.

Example: $\mathrm{CH_3COOH} + \mathrm{NH_3} \rightleftharpoons \mathrm{CH_3COO^-} + \mathrm{NH_4^+}$

pH and Strength

• pH measures the concentration of H+ ions: $\mathrm{pH} = -\log[\mathrm{H^+}]$

• High [H+]: low pH (acidic); low [H+]: high pH (basic).

Strong vs. Weak Acids

• Strong acids dissociate completely in water.

• Weak acids establish an equilibrium with their ions.

Example: $\mathrm{HCl} \rightarrow \mathrm{H^+} + \mathrm{Cl^-}$ (strong acid)

$\mathrm{CH_3COOH} \rightleftharpoons \mathrm{CH_3COO^-} + \mathrm{H^+}$ (weak acid)

Acid Dissociation Constant ($K_a$) and $pK_a$

• $K_a$ measures acid strength: $K_a = \frac{[\mathrm{H^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$

• Larger $K_a$: stronger acid; smaller $K_a$: weaker acid.

• $pK_a = -\log K_a$; lower $pK_a$ means stronger acid.

Buffers

Definition and Function

• A buffer is a mixture of a weak acid and its conjugate base.

• Buffers resist large changes in pH when small amounts of acid or base are added.

Buffer Action and Le Châtelier’s Principle

• Adding strong acid to a buffer shifts equilibrium to reduce the increase in [H+].

• pH change is minimized compared to pure water.

Calculating Buffer pH: Henderson-Hasselbalch Equation

• The pH of a buffer can be calculated using:

$\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[\text{base}]}{[\text{acid}]} \right)$

• Example: For 0.200 M acetic acid and 0.300 M sodium acetate, $K_a = 1.8 \times 10^{-5}$:

$\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{0.300}{0.200} \right)$

Additional info: The Henderson-Hasselbalch equation is derived from the equilibrium expression for weak acids and is widely used in biochemistry and medicine.