CHEM 1030 Exam 1 Master Study Guide

What should I study for lab safety?

All safety regulations and proper lab techniques from the first-week handout. Be able to recognize the safe response to situations covered in class.

Does this guide replace the safety handout?

No. Review the actual handout and class demonstrations.

What is chemistry?

The study of matter, its properties, and the changes it undergoes.

What are the steps of the scientific method?

Observe, ask a question, form a testable hypothesis, experiment, analyze results, and draw a conclusion.

What is matter?

Anything that has mass and occupies space.

Element vs. compound

An element contains one type of atom. A compound contains two or more elements chemically combined in fixed proportions.

Substance vs. mixture

A substance has a fixed composition; a mixture is a physical combination with variable composition.

Homogeneous vs. heterogeneous mixture

Homogeneous: uniform throughout. Heterogeneous: nonuniform throughout.

What are the three common states of matter?

Solid: definite shape and volume. Liquid: definite volume, takes container shape. Gas: no definite shape or volume.

Physical property

Observed without changing chemical identity; examples include color, density, melting point, and boiling point.

Chemical property

Describes the ability to undergo chemical change, such as flammability or reactivity.

Physical vs. chemical change

Physical change alters form/state but not identity. Chemical change forms different substances (rusting or burning).

Common SI units used in introductory chemistry

Mass: kilogram (kg); length: meter (m); time: second (s). Chemistry often uses grams, centimeters, and seconds.

Density formula

D = m/V. Rearranged: m = D×V; V = m/D.

Accuracy vs. precision

Accuracy is closeness to the accepted value. Precision is closeness of repeated measurements to one another.

Multiplication/division sig-fig rule

Round to the fewest significant figures in the measured values.

Addition/subtraction sig-fig rule

Round to the least precise decimal place.

Common dimensional-analysis errors

Flipping a conversion factor, failing to cancel units, dropping units, or forgetting to square the factor for area/volume.

Three subatomic particles, charges, and locations

Proton: +1, nucleus. Neutron: 0, nucleus. Electron: -1, outside nucleus.

Atomic number

Number of protons; it identifies the element.

Mass number

Protons + neutrons.

How do you calculate neutrons?

Neutrons = mass number - atomic number.

How do you find electrons in ions?

Positive charge means electrons were lost (electrons = protons - charge). Negative charge means electrons were gained (add magnitude of charge).

What are isotopes?

Atoms of the same element with the same proton number but different neutron numbers/mass numbers.

Group vs. period

Group = vertical column; period = horizontal row.

Common group names

Group 1: alkali metals; Group 2: alkaline earth metals; Group 17: halogens; Group 18: noble gases.

Ionic compound

Made of cations and anions; the overall formula is electrically neutral.

Molecular compound

A covalent compound formed from nonmetals; its formula represents molecules.

Central rule for writing ionic formulas

Choose subscripts so total positive and negative charges sum to zero, then reduce to the lowest whole-number ratio.

Subscript vs. coefficient

A subscript is part of a formula and changes particle composition. A coefficient multiplies the entire formula.

Formula unit vs. molecule

Formula unit is the simplest ratio of ions in an ionic compound; molecule is a discrete covalently bonded unit.

What does a chemical equation show?

Reactants on the left, products on the right; coefficients show relative amounts. Balance each element's atom count.

Common hydrocarbon combustion pattern

Hydrocarbon + O₂ -> CO₂ + H₂O. Balance carbon, then hydrogen, then oxygen.

What is one mole?

6.022 × 10²³ particles (Avogadro's number).

How do you calculate molar mass?

Add the atomic masses of all atoms in the formula, multiplying each by its subscript (g/mol).

Molecular mass vs. molar mass

Molecular mass is mass of one molecule in u; molar mass is mass of one mole in g/mol.

Grams to moles conversion

Moles = grams ÷ molar mass.

Moles to grams conversion

Grams = moles × molar mass.

Percent composition by mass

(Mass of element in 1 mol compound ÷ molar mass of compound) × 100%.

Empirical vs. molecular formula

Empirical = simplest whole-number atom ratio. Molecular = actual number of each atom in a molecule.

Steps for empirical formula from percentages

Assume 100 g; change percentages to grams; convert grams to moles; divide by smallest mole value; multiply to whole numbers if needed.

Stoichiometry

Using a balanced equation to calculate amounts of reactants and products.

Limiting vs. excess reactant

Limiting reactant is consumed first and limits product; excess reactant remains.

Theoretical yield

Maximum product predicted from the limiting reactant.

Actual yield

Amount of product actually obtained.

Percent yield formula

(Actual yield ÷ theoretical yield) × 100%.