Chemical Equations & Bonding

Cheatsheet Content

1. Chemical Equations Definition: A symbolic representation of a chemical reaction, showing reactants and products. Reactants: Substances that undergo change, on the left side of the arrow. Products: Substances formed as a result of the reaction, on the right side of the arrow. Arrow ($\rightarrow$): "Yields" or "produces". States of Matter: $(s)$ - solid $(l)$ - liquid $(g)$ - gas $(aq)$ - aqueous (dissolved in water) Balancing Equations Law of Conservation of Mass: Atoms are neither created nor destroyed in a chemical reaction. Number of atoms for each element must be equal on both sides. Use coefficients (numbers in front of formulas) to balance. Steps: Write the unbalanced equation. Count atoms of each element on both sides. Balance elements one by one (often metals first, then non-metals, then H and O last). Check all elements. Example: $\text{CH}_4 + \text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O}$ (unbalanced) Balanced: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$ 2. Types of Chemical Reactions Synthesis (Combination): $A + B \rightarrow AB$ Example: $2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}$ Decomposition: $AB \rightarrow A + B$ Example: $2\text{H}_2\text{O} \rightarrow 2\text{H}_2 + \text{O}_2$ Single Displacement (Single Replacement): $A + BC \rightarrow AC + B$ $A$ must be more reactive than $B$. Example: $\text{Zn} + \text{CuCl}_2 \rightarrow \text{ZnCl}_2 + \text{Cu}$ Double Displacement (Double Replacement): $AB + CD \rightarrow AD + CB$ Often forms a precipitate, gas, or water. Example: $\text{AgNO}_3(aq) + \text{NaCl}(aq) \rightarrow \text{AgCl}(s) + \text{NaNO}_3(aq)$ Combustion: Reaction with oxygen, often producing heat and light. Hydrocarbon combustion: $\text{hydrocarbon} + \text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O}$ Example: $\text{C}_3\text{H}_8 + 5\text{O}_2 \rightarrow 3\text{CO}_2 + 4\text{H}_2\text{O}$ Acid-Base (Neutralization): Acid + Base $\rightarrow$ Salt + Water Example: $\text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$ 3. Chemical Bonding Definition: The force that holds atoms together to form compounds. Atoms bond to achieve a more stable electron configuration (usually a full valence shell - octet rule). Ionic Bonding Formation: Transfer of electrons between a metal (loses electrons to form cation) and a non-metal (gains electrons to form anion). Electronegativity Difference: Large (typically $> 1.7$). Characteristics: Form crystal lattices. High melting/boiling points. Conduct electricity when molten or dissolved in water. Brittle. Example: $\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl}$ Covalent Bonding Formation: Sharing of electrons between two non-metals. Electronegativity Difference: Small or zero (typically $ Types: Nonpolar Covalent: Equal sharing of electrons (electronegativity difference $ Polar Covalent: Unequal sharing of electrons, creating partial positive ($\delta^+$) and partial negative ($\delta^-$) charges (electronegativity difference $0.4 - 1.7$). Example: $\text{H}_2\text{O}$ Characteristics: Form discrete molecules. Lower melting/boiling points than ionic compounds. Poor electrical conductors. Lewis Structures: Diagrams showing valence electrons as dots and shared pairs as lines. Metallic Bonding Formation: Between metal atoms, where valence electrons are delocalized and form a "sea of electrons" shared among all atoms. Characteristics: High electrical and thermal conductivity. Malleable (can be hammered into sheets) and ductile (can be drawn into wires). Lustrous (shiny). High melting/boiling points. 4. Intermolecular Forces (IMFs) Forces of attraction between molecules (weaker than intramolecular bonds). Influence physical properties like melting point, boiling point, viscosity. Types of IMFs (increasing strength) London Dispersion Forces (LDF): Present in all molecules (polar and nonpolar). Caused by temporary, instantaneous dipoles due to electron movement. Strength increases with molecular size/number of electrons. Dipole-Dipole Forces: Present in polar molecules. Attraction between permanent dipoles of adjacent molecules. Hydrogen Bonding: A special, strong type of dipole-dipole force. Occurs when H is bonded to a highly electronegative atom (N, O, or F). Responsible for high boiling point of water. 5. Molecular Geometry (VSEPR Theory) Valence Shell Electron Pair Repulsion Theory: Electron domains (lone pairs and bonding pairs) repel each other and arrange to minimize repulsion. Determines the 3D shape of molecules. Electron Domains Electron Geometry Bonding Pairs Lone Pairs Molecular Geometry Example 2 Linear 2 0 Linear $\text{CO}_2$ 3 Trigonal Planar 3 0 Trigonal Planar $\text{BF}_3$ 3 Trigonal Planar 2 1 Bent $\text{SO}_2$ 4 Tetrahedral 4 0 Tetrahedral $\text{CH}_4$ 4 Tetrahedral 3 1 Trigonal Pyramidal $\text{NH}_3$ 4 Tetrahedral 2 2 Bent $\text{H}_2\text{O}$ 5 Trigonal Bipyramidal 5 0 Trigonal Bipyramidal $\text{PCl}_5$ 5 Trigonal Bipyramidal 4 1 Seesaw $\text{SF}_4$ 5 Trigonal Bipyramidal 3 2 T-shaped $\text{ClF}_3$ 5 Trigonal Bipyramidal 2 3 Linear $\text{XeF}_2$ 6 Octahedral 6 0 Octahedral $\text{SF}_6$ 6 Octahedral 5 1 Square Pyramidal $\text{BrF}_5$ 6 Octahedral 4 2 Square Planar $\text{XeF}_4$