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Chapter 1 — Chemistry skills and scientific thinking

• Chemistry studies matter and the changes matter undergoes.
• A hypothesis is testable; a theory is a well-supported explanation.
• Independent variable = what is changed; dependent variable = what is measured.
• Qualitative observations describe; quantitative observations measure.
• Every measurement needs a number and a unit.
• Exact values come from counting or definitions; measured values have uncertainty.

Chapter 2 — Measurements, conversions, and significant figures

• Metric prefixes: kilo = 10³, centi = 10⁻², milli = 10⁻³, micro = 10⁻⁶.
• Use dimensional analysis: given × conversion factor = needed unit.
• Multiplication/division: round to the fewest significant figures.
• Addition/subtraction: round to the fewest decimal places.
• Density = mass / volume; mass = density × volume; volume = mass / density.
• K = °C + 273.15 and °C = (°F − 32) / 1.8.

Chapter 3 — Matter, energy, and heat

• Pure substances are elements or compounds; mixtures can be homogeneous or heterogeneous.
• Physical changes do not form new substances; chemical changes do.
• Solids have definite shape and volume; liquids have definite volume; gases have neither.
• Exothermic releases heat; endothermic absorbs heat.
• q = m c ΔT for heat transfer calculations.
• Mass is conserved in a closed system.

Chapter 4 — Atoms, elements, isotopes, and ions

• Atomic number = number of protons and defines the element.
• Mass number = protons + neutrons.
• Neutral atoms have protons = electrons.
• Isotopes have the same protons but different neutrons.
• Cations are positive and usually form by losing electrons.
• Anions are negative and usually form by gaining electrons.

Chapter 5 — Valence electrons, periodic trends, radiation, and nuclear chemistry

• Main-group number helps predict valence electrons.
• Group 1 has 1 valence electron; Group 2 has 2; Group 17 has 7; Group 18 has full shells.
• Metals tend to lose electrons; nonmetals tend to gain electrons.
• Atomic radius decreases across a period and increases down a group.
• Ionization energy and electronegativity generally increase toward the upper right.
• Lewis symbols show valence electrons as dots.
• Isotopes of the same element have the same atomic number but different mass numbers because they have different neutron counts.
• Mass number = protons + neutrons; atomic number = protons.
• Radioactive nuclei emit radiation to become more stable.
• Alpha particles have mass number 4 and charge 2+; beta particles have mass number 0 and charge 1−; positrons have mass number 0 and charge 1+; gamma rays have mass number 0 and charge 0.
• In nuclear equations, total mass numbers and total atomic numbers must balance on both sides.
• Alpha decay lowers mass number by 4 and atomic number by 2; beta decay leaves mass number unchanged and increases atomic number by 1; positron emission leaves mass number unchanged and decreases atomic number by 1; gamma emission changes neither.
• Half-life is the time needed for half of a radioactive sample to decay.
• Radiation dose conversions include 1 mCi = 1000 microCi, 1 rad = 1000 mrad, 1 Gy = 100 rad, and for gamma radiation 1 rad is approximately 1 rem.

Chapter 6A — Ions and the Octet Rule

• Atoms form ions or bonds to reach a more stable valence electron arrangement.
• Main-group metals usually lose electrons and become cations.
• Main-group nonmetals usually gain electrons and become anions.
• Group 1 forms 1+ ions, Group 2 forms 2+ ions, Group 16 often forms 2− ions, and Group 17 often forms 1− ions.
• A Lewis dot symbol shows valence electrons around an element symbol.

Chapter 6B — Ionic Compounds

• Ionic compounds form from cations and anions and must have zero net charge.
• Write the cation first and the anion second in an ionic formula.
• Use subscripts to show the smallest whole-number ratio that balances charges.
• Binary ionic nonmetals end in -ide, such as chloride, oxide, and nitride.
• Do not use molecular prefixes like di- or tri- when naming ordinary ionic compounds.

Chapter 6C — Transition Metals and Roman Numerals

• Many transition metals can form more than one positive charge.
• A Roman numeral in the name shows the metal ion charge, such as iron(III) for Fe³⁺.
• Use the known anion charge and the neutral formula to calculate the transition metal charge.
• Metals with fixed common charges, such as Group 1, Group 2, aluminum, zinc, silver, and cadmium, usually do not need Roman numerals.

Chapter 6D — Polyatomic Ions

• Polyatomic ions are charged groups of covalently bonded atoms that act as one ion.
• Memorize common ions such as ammonium, hydroxide, nitrate, sulfate, carbonate, phosphate, chlorate, and acetate.
• Use parentheses when a formula needs more than one copy of a polyatomic ion.
• Do not change the atoms inside a polyatomic ion when balancing charges.
• Many oxyanion pairs follow -ate for more oxygen and -ite for fewer oxygen atoms.

Chapter 6E — Covalent Compounds

• Covalent compounds usually form between nonmetals that share electrons.
• Molecular compounds use prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca.
• The first element keeps its element name and usually skips mono- when there is only one atom.
• The second element uses a prefix and ends in -ide.
• Molecular formulas show actual atom counts, not charge-balanced ion ratios.

Chapter 6F — Bond Polarity and Attractive Forces

• Electronegativity differences help predict nonpolar covalent, polar covalent, or ionic bonding.
• Polar bonds have unequal electron sharing and partial charges.
• All particles have dispersion forces; larger electron clouds usually create stronger dispersion forces.
• Polar molecules can have dipole-dipole attractions.
• Hydrogen bonding occurs when H is bonded to N, O, or F and is attracted to a lone pair on N, O, or F nearby.