General Chemistry: Atoms, Bonding, Lewis Structures, and Molecular Geometry

Valence Electrons

These are the outermost electrons of an atom involved in chemical bonding.

Atomic Number

This number represents the number of protons in the nucleus of an atom.

Mass Number

The mass number is the sum of protons and neutrons in an atom's nucleus.

Atomic Weight

This is the weighted average mass of all naturally occurring isotopes of an element.

Calculating Neutrons

To find the number of neutrons subtract the atomic number from the mass number (for example Selenium has a mass of 79 and atomic number of 34 yielding 45 neutrons).

Electronegativity

This is a measure of an atom's ability to attract shared electrons in a chemical bond.

Lewis Dot Structures

These are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist.

Unit Conversions and Sigfigs

These are foundational mathematical skills required to properly calculate and format stoichiometry and molar mass answers.

Chemical Reaction

A process where one or more starting chemicals are converted into one or more chemical products.

Reactants

The starting materials or chemicals in a chemical reaction.

Products

The ending materials or chemicals produced in a chemical reaction.

Octet Rule

A rule of thumb stating that main group elements tend to gain or lose electrons to achieve a full outer shell of eight valence electrons.

Anion

An ion with a negative charge formed when an atom gains one or more electrons to fill its outermost shell.

Cation

An ion with a positive charge formed when an atom loses one or more electrons to empty its outermost shell.

Ionic Bond

A chemical bond resulting from the attraction between positive and negative ions usually forming when the electronegativity difference is greater than 1.8.

Covalent Bond

A chemical bond resulting from the sharing of electrons between two atoms generally occurring when the electronegativity difference is less than 0.5.

Polar Covalent Bond

A bond where electrons are not shared equally between two atoms typically occurring when the electronegativity difference is between 0.5 and 1.8.

Nonpolar Covalent Bond

A bond where electrons are shared equally between two atoms because their electronegativity difference is very small.

Monatomic Cations

These metals form only one type of cation and are named simply by adding ion after the element name.

Transition Metal Cations

These metals can form more than one type of cation so their charge must be denoted using Roman numerals in parentheses.

Monatomic Anions

These negative ions are named by taking the stem of the element name and adding the suffix -ide.

Polyatomic Ions

These are tightly bound groups of atoms that carry an overall charge and have specific common names you must memorize.

Nitrite Ion

This is a specific polyatomic anion with the formula NO2- carrying a negative one charge.

Hydrogen Ion

This is a simple cation with the formula H+ representing a proton.

Fluoride Ion

This is a monatomic anion with the formula F- formed when fluorine gains one electron.

Dihydrogen Phosphate Ion

This is a polyatomic anion with the formula H2PO4- carrying a negative one charge.

Sulfate Ion

This is a polyatomic anion with the formula SO42- carrying a negative two charge.

Naming Ionic Compounds

Write the name of the positively charged cation first followed by the name of the negatively charged anion.

Naming Binary Covalent Compounds

Name the less electronegative element first then name the more electronegative element ending in -ide and use prefixes like di- and tri-.

Bonding Electrons

These are the valence electrons that are actively shared between atoms to form a covalent bond.

Nonbonding Electrons

These are valence electrons that reside on a single atom and are not involved in forming covalent bonds.

Single Bond

A covalent bond formed by sharing one pair of electrons represented by a single line in a Lewis structure.

Double Bond

A covalent bond formed by sharing two pairs of electrons represented by two lines between the bonded atoms.

Triple Bond

A covalent bond formed by sharing three pairs of electrons represented by three lines between the bonded atoms.

Drawing Lewis Structures Step 1

First determine the total number of valence electrons in the molecule by adding up the valence electrons contributed by each atom.

Drawing Lewis Structures Step 2

Determine the connectivity of the atoms and connect bonded atoms with single bonds noting that carbon nitrogen oxygen and halogens need eight electrons.

Exceptions to the Octet Rule

Period 3 elements can have expanded octets of 10 or 12 electrons because their 3s electrons can be promoted to new orbitals.

Resonance

This occurs when electrons are shared among multiple bonds making it impossible to draw a single static structure.

Bicarbonate Resonance

The bicarbonate ion HCO3- exhibits resonance because its double bond and lone pairs can be distributed in multiple valid structural configurations.

Contributing Structures

These are the individual Lewis structures that make up a resonance hybrid differing only in the distribution of their valence electrons.

Resonance Hybrid

A molecule or ion described as a composite or hybrid of several contributing resonance structures.

VSEPR Model

The valence-shell electron-pair repulsion model predicts 3D bond angles and shapes based on the idea that electron regions repel each other.

Tetrahedral Shape

A molecular shape formed when 4 identical atoms or groups are attached to a central atom resulting in bond angles of 109.5 degrees.

Pyramidal Shape

A shape similar to tetrahedral but with one lone pair of electrons instead of a bonded atom resulting in a compressed bond angle of 107.3 degrees.

Bent Shape

A modified tetrahedral shape with two lone pairs of electrons resulting in a compressed bond angle of 104.5 degrees.

Trigonal Planar Shape

A molecular shape involving 3 regions of electron density around a central atom resulting in bond angles of 120 degrees.

Linear Shape

A molecular shape involving 2 regions of electron density around a central atom resulting in bond angles of 180 degrees.

Methane Geometry

Methane (CH4) has four bonded hydrogen atoms and zero lone pairs around the central carbon making it a perfect tetrahedral shape with 109.5-degree angles.

Ammonia Geometry

Ammonia (NH3) has three bonded hydrogen atoms and one lone pair around the central nitrogen making its geometry pyramidal.

Molecular Polarity

A molecule is polar if it contains polar bonds that are distributed unevenly and do not cancel each other out.

Dipole Cancellation

If a molecule is highly symmetrical the polar bonds cancel each other out completely making the overall molecule nonpolar.

Formula Weight

The sum of the atomic weights in amu of all atoms in a compound's formula used for both ionic and covalent compounds.

Molecular Weight

The sum of the atomic weights in amu of all atoms in molecules held together strictly by covalent bonds.

The Mole

The number of atoms in exactly 12.000 grams of carbon-12 which equals Avogadro's number.

Avogadro's Number

The constant 6.02 x 10^23 which represents the number of individual atoms or molecules in one mole of a substance.

Molar Mass

The mass of one mole of a substance expressed in grams per mole calculated by adding the formula weights.

Converting Grams to Moles

When solving stoichiometry problems divide the given grams by the compound's total molar mass.

Converting Moles to Molecules

To find the number of individual molecules or atoms multiply the given number of moles by Avogadro's number.

Balancing Chemical Equations

Ensure that the exact atom count for each element is identical on both the reactant and product sides by adjusting coefficients.

Stoichiometry Step 1

Always begin by writing and verifying the perfectly balanced chemical equation to establish the correct mole ratios.

Stoichiometry Step 2

Convert the given mass of the starting materials into moles using their specific molar mass from the periodic table.

Stoichiometry Step 3

Multiply your starting moles by the ratio of the coefficients from the balanced equation to solve for the moles of product.

Stoichiometry Step 4

Convert your calculated product moles back into a final gram amount by multiplying by the product's molar mass.

Limiting Reagent

The reactant that mathematically runs out first and dictates the maximum possible amount of product that can be formed.

Actual Yield

The measurable physical mass of a product that is literally formed and recovered during a real chemical reaction.

Theoretical Yield

The maximum calculated mass of a product that mathematically should form according to the stoichiometry of the limiting reagent.

Percent Yield

The ratio of the actual yield divided by the theoretical yield multiplied by 100 to show reaction efficiency.

Aqueous Solution

A specific type of solution in which the solvent dissolving the chemical compounds is liquid water.

Soluble Ionic Compounds

Ionic compounds that completely dissociate into individual positive and negative free-floating ions when placed in water.

Insoluble Ionic Compounds

Ionic compounds that do not dissolve in water and instead stay strongly bound together as a solid precipitate.

Complete Ionic Equation

An equation that shows all soluble strong electrolytes completely dissociated into their individual charged ions.

Spectator Ions

Ions that remain dissolved identically on both the reactant and product sides without participating in the actual chemical change.

Net Ionic Equation

A simplified equation showing only the specific aqueous ions that chemically combine to form the solid precipitate.

Oxidation

A chemical process defined by the loss of electrons, an increase in charge, the loss of hydrogen, or the gain of oxygen.

Reduction

A chemical process defined by the gain of electrons, a decrease in charge, the gain of hydrogen, or the loss of oxygen.

Redox Reaction

An oxidation-reduction reaction where electrons are fundamentally transferred from one participating reactant to another.

Oxidizing Agent

The entity that accepts electrons causing another substance to be oxidized while it simultaneously becomes reduced itself.

Reducing Agent

The entity that donates electrons causing another substance to be reduced while it simultaneously becomes oxidized itself.

Exothermic Reaction

A chemical reaction that actively produces and gives off heat energy to its immediate surroundings.

Endothermic Reaction

A chemical reaction that requires and actively absorbs heat energy from its immediate surroundings to proceed.

Example Problem: If you burn 1.20 moles of butane (C4H10) in oxygen how many moles of CO2 will be generated?

To solve this, first write the balanced equation: 2C4H10 + 13O2 -> 8CO2 + 10H2O. Next, find the mole ratio between the reactant and product, which is 2 moles of butane to 8 moles of CO2. Finally, multiply your starting amount by this ratio: 1.20 moles of C4H10 * (8 moles CO2 / 2 moles C4H10) to get 4.80 moles of CO2.

Example Problem: How many moles of methane (CH4) are in 32.0 g of CH4?

First calculate the molar mass of CH4 by adding the atomic weights of one Carbon (12.0 g) and four Hydrogens (4.0 g) to get 16.0 g/mol. Next set up your conversion factor dividing the given grams by the molar mass. Divide 32.0 g by 16.0 g/mol to find that you have exactly 2.00 moles of CH4.

Example Problem: If a reaction forming ethanol has a theoretical yield of 50.5 g and an actual yield of 46.8 g what is the percent yield?

Percent yield compares the amount you actually produced to the maximum amount you mathematically could have produced. To solve use the formula: (Actual Yield / Theoretical Yield) * 100%. Divide 46.8 g by 50.5 g then multiply by 100 to get a percent yield of 92.7%.

Example Problem: What is the correct chemical formula for Calcium phosphate?

First identify the charges of the individual ions: Calcium is in Group 2A so it forms a Ca2+ cation and phosphate is a polyatomic anion with the formula PO4 3-. Next find the least common denominator for the charges which is 6. To make the compound neutral you need three Ca2+ ions and two PO4 3- ions giving the final formula Ca3(PO4)2.

Example Problem: What is the net ionic equation when mixing NaCl(aq) and AgNO3(aq)?

First recognize that they form the soluble ions Na+ Cl- Ag+ and NO3- which combine to form the insoluble precipitate AgCl(s). The Na+ and NO3- ions remain dissolved and unchanged on both sides making them spectator ions. Cross out the spectator ions to leave only the ions forming the solid: Ag+(aq) + Cl-(aq) -> AgCl(s).

Example Problem: In the reaction Pb(s) + 2Ag+(aq) -> Pb2+(aq) + 2Ag(s) what is the oxidizing agent and what is the reducing agent?

First analyze the change in charges to determine what is oxidized and reduced. Lead (Pb) goes from 0 to +2 meaning it lost electrons (oxidized) while Silver (Ag+) goes from +1 to 0 meaning it gained electrons (reduced). The substance that is oxidized acts as the reducing agent (Pb) and the substance that is reduced acts as the oxidizing agent (Ag+).

Example Problem: How do you draw the Lewis structure for Formaldehyde (CH2O)?

First calculate the total valence electrons: 4 from C 2 from the two H's and 6 from O totaling 12 electrons. Next arrange the atoms with Carbon in the center connecting it to the Hydrogens and Oxygen with single bonds which uses 6 electrons. Finally to satisfy the octet rule for Carbon and Oxygen create a double bond between C and O and place the remaining 4 electrons as two nonbonding pairs on the Oxygen.

Example Problem: Based on VSEPR what is the shape and polarity of Water (H2O)?

Draw the Lewis structure to see Oxygen in the center bonded to two hydrogens, with two unshared electron pairs on the Oxygen. These four regions of electron density create a bent molecular shape with compressed 104.5-degree bond angles. Because Oxygen is highly electronegative and the bent shape is asymmetrical, the dipoles do not cancel, making water a polar molecule.

Example Problem: How do you determine if the bond between Carbon (C) and Oxygen (O) is ionic polar covalent or nonpolar covalent?

First locate the electronegativity values for each element on the provided scale: Carbon is 2.5 and Oxygen is 3.5. Next calculate the absolute difference between these two values which is 1.0. Because 1.0 falls between 0.5 and 1.9 this is classified as a polar covalent bond.

Example Problem: How many grams of H2O are needed to fully react with 37.05 g of Ca(OH)2 in a balanced equation?

First ensure you have a properly balanced chemical equation to establish your mole-to-mole ratio between Ca(OH)2 and H2O. Next divide the 37.05 g by the molar mass of Ca(OH)2 to find your starting moles. Finally multiply those starting moles by the balanced mole ratio and then multiply by the molar mass of H2O (18.0 g/mol) to get your final grams.

CaCl2 Equation Step 1: Balancing

How do you balance Ca(OH)2 + HCl -> CaCl2 + H2O? First count your atoms on both sides. Place a 2 in front of the HCl to balance the chlorines then place a 2 in front of the H2O to balance the hydrogen and oxygen resulting in 1 Ca(OH)2 + 2 HCl -> 1 CaCl2 + 2 H2O.

CaCl2 Equation Step 2: Reactant Moles

How do you find the starting moles of 37.05 g of Ca(OH)2? Calculate the molar mass of Ca(OH)2 (40.08 + 2(16.00) + 2(1.01) = 74.10 g/mol) then divide the starting mass of 37.05 g by 74.10 g/mol to get 0.5000 moles.

CaCl2 Equation Step 3: Product Grams

How do you find the grams of CaCl2 produced from 0.5000 moles of Ca(OH)2? Multiply the 0.5000 moles by the 1:1 balanced equation ratio to get 0.5000 moles of CaCl2 then multiply by CaCl2's molar mass (110.98 g/mol) to get 55.49 g of CaCl2.

CaCl2 Equation Step 4: Product Molecules

How do you find the molecules of CaCl2 produced? Multiply the 0.5000 moles of CaCl2 by Avogadro's number (6.022 x 10^23) to get 3.011 x 10^23 molecules of CaCl2.