Chemistry Unit 1

Repeated lessons 1-3 but all from the study guide Format Atomic Structure — subatomic particles, Thomson, Rutherford, Chadwick, Bohr, quantized energy Ions & Octet Rule — cation/anion formation, group charges, polyatomic ions Isotopes & Average Atomic Mass — formula, radioisotopes, real-world applications Periodic Table Structure — groups/periods, Alkali Metals, Halogens, Noble Gases Periodic Trends — Zeff, shielding, AR, EN, IE, EA, reactivity, melting point Lewis Structures & Bonding — drawing steps, bond types, coordinate covalent bonds Nomenclature — ionic, multivalent, covalent, acids, hydrates, oxyacid salts Polarity — ΔEN benchmarks, dipole moments, determining molecular polarity IMFs — LDF, dipole-dipole, H-bonding, bulk properties, like dissolves like Quick Reference — all common polyatomic ions

What are the 3 subatomic particles?

Proton (+1, in nucleus), Neutron (0, in nucleus), Electron (–1, orbits nucleus in energy levels)

What is the net charge of a neutral atom?

Zero — the number of protons equals the number of electrons.

Where are protons and neutrons found?

In the dense nucleus at the centre of the atom.

Where are electrons found?

In fixed energy levels (orbits/shells) surrounding the nucleus — they occupy a probability density cloud.

How do you find the number of neutrons?

Round the atomic mass to the nearest whole number, then subtract the atomic number (# of protons).

Who discovered the electron, and with what model?

J.J. Thomson — he proposed the 'Raisin Bun' model: a positive sphere with electrons embedded.

What did Rutherford's gold foil experiment show?

The atom is mostly empty space with a dense, positive nucleus. Most alpha particles passed through; ~1/20,000 bounced back 180°.

Who discovered the neutron and why was it important?

James Chadwick (1932) — the neutron explained why an all-positive nucleus doesn't repel itself and accounted for the atom's extra mass.

What is the Bohr model?

Electrons occupy fixed, circular energy levels (stationary states). They emit/absorb quantized energy when jumping between levels.

What does 'quantized energy' mean?

Energy exists only in discrete amounts (chunks called quanta) — like steps on a staircase, not a continuous slope.

Why are elements ordered by atomic number, not atomic mass?

Some pairs (Ar/K, Co/Ni, Te/I) are out of order by mass. Ordering by proton count (atomic number) corrects this. Mendeleev didn't know about subatomic particles.

What is an ion?

An atom that has gained or lost electrons, giving it a net electrical charge.

What is a cation?

A positively charged ion — formed when a metal LOSES electrons. (e.g., Na⁺, Ca²⁺)

What is an anion?

A negatively charged ion — formed when a nonmetal GAINS electrons. (e.g., Cl⁻, O²⁻)

What is the Octet Rule?

Atoms gain, lose, or share electrons to achieve 8 valence electrons in their outer shell (like a noble gas). Exception: H uses the Duet Rule (2 e⁻).

What ionic charges do Groups 1–3 metals form?

Group 1 → 1+, Group 2 → 2+, Group 3 → 3+

What ionic charges do Groups 15–17 nonmetals form?

Group 15 → 3–, Group 16 → 2–, Group 17 → 1–

What is a polyatomic ion?

A group of covalently bonded atoms that carries an overall ionic charge (e.g., NO₃⁻ nitrate, SO₄²⁻ sulfate, NH₄⁺ ammonium).

What are isotopes?

Atoms of the same element (same protons) with different numbers of neutrons — so they have different atomic masses.

Do isotopes behave the same chemically?

Yes — same electron configuration = same chemical behaviour. But some isotopes are radioactive due to nuclear instability.

What is average atomic mass?

A weighted average of all naturally occurring isotopes of an element, accounting for each isotope's mass AND percent abundance.

Formula for average atomic mass?

AMavg = (M₁)(abundance₁) + (M₂)(abundance₂) + … where M = mass in amu, abundance = decimal fraction.

Example: Carbon's average atomic mass?

AMavg = (12)(0.9889) + (13)(0.01108) + (14)(1.0×10⁻¹⁴) ≈ 12.011 amu

What is C-12 used as and why?

The standard for the atomic mass unit (amu): 1 amu = 1/12 the mass of a C-12 atom. Chosen for high abundance, stability, and solid state.

What are radioisotopes?

Unstable isotopes that decay, releasing alpha/beta particles or gamma rays to become more stable nuclei.

Name 5 isotope applications.

C-14 → radiocarbon dating; Co-60 → cancer radiation therapy; Tc-99m → CT scans; U-235 → nuclear reactors; H-3 (Tritium) → glowing watch faces/optics

What is a period on the periodic table?

A horizontal ROW. Elements in the same period have the same number of energy levels.

What is a group (family) on the periodic table?

A vertical COLUMN. Elements in the same group have the same number of valence electrons and similar properties.

Properties of Alkali Metals (Group 1)?

1 valence e⁻; very soft, shiny, low melting point; react violently with water to produce H₂ gas and a basic solution; most reactive metals.

1 valence e⁻; very soft, shiny, low melting point; react violently with water to produce H₂ gas and a basic solution; most reactive metals.

2 valence e⁻; harder than alkali metals; react with acids and O₂; oxides produce basic solutions with water.

Properties of Halogens (Group 17)?

7 valence e⁻; very reactive (gain 1 e⁻); F₂ and Cl₂ are gases, Br₂ is liquid, I₂ is solid at room temp; fluorine is the most reactive element.

Properties of Noble Gases (Group 18)?

Full valence shells; colourless, odourless gases; chemically inert (no electron affinity). Helium is stable with 2 e⁻ (duet rule).

How many valence electrons do elements in Group 14 have?

4 valence electrons (e.g., C, Si)

What is Effective Nuclear Charge (Zeff)?

The net positive charge experienced by an outer electron — the 'pull' the nucleus exerts on valence electrons, after accounting for shielding.

What is the shielding effect?

Inner-shell electrons block outer electrons from experiencing the full attraction of the nucleus, reducing Zeff.

How does Zeff change across a period (left → right)?

INCREASES — more protons added, same number of energy levels, so shielding stays constant → stronger nuclear pull on valence electrons.

How does Zeff change down a group?

DECREASES (effectively) — more energy levels added → greater shielding → outer electrons feel less nuclear pull.

What is atomic radius (AR)?

The distance from the nucleus to the outermost electrons — effectively the 'size' of the atom.

How does AR change across a period?

DECREASES — more protons increase Zeff, pulling electrons closer. Same number of energy levels, so no extra shielding.

How does AR change down a group?

INCREASES — more energy levels are added, pushing outer electrons farther from the nucleus.

What is electronegativity (EN)?

The ability of an atom to attract a shared bonding pair of electrons toward itself (measured on the Pauling scale).

How does EN change across a period?

INCREASES — higher Zeff pulls bonding electrons more strongly.

How does EN change down a group?

DECREASES — greater shielding and larger radius reduce the nucleus's ability to attract electrons.

Which element has the highest electronegativity?

Fluorine (F) — top-right of the periodic table (excluding noble gases).

What is ionization energy (IE)?

The energy required to remove the least tightly held electron from a gaseous atom: A(g) + energy → A⁺(g) + e⁻

How does IE change across a period?

INCREASES — higher Zeff means electrons are held more tightly, requiring more energy to remove.

How does IE change down a group?

DECREASES — greater AR and shielding reduce Zeff, so less energy is needed to remove an outer electron.

What is electron affinity (EA)?

The energy change (released or absorbed) when a neutral gaseous atom gains a free electron: A(g) + e⁻ → A⁻(g). Expressed as a negative value.

How does EA change across a period and down a group?

Across a period: EA increases (more negative). Down a group: EA decreases (less negative). Noble gases have no EA.

Reactivity trend for metals?

Increases DOWN a group (easier to lose outer electrons) and DECREASES across a period (higher Zeff holds electrons tighter).

Reactivity trend for nonmetals?

Increases UP a group (smaller atoms attract electrons more strongly) and INCREASES across a period (higher Zeff).

Melting point trend for metals vs. nonmetals?

Metals: generally decreases down a group. Nonmetals: generally increases down a group.

Steps to draw a Lewis structure (overview)?

1) Arrange atoms (least electronegative in centre; H always peripheral). 2) Count total valence e⁻. 3) Place single bonds. 4) Complete octets on outer atoms. 5) Place remaining e⁻ on central atom. 6) If central atom incomplete, convert lone pairs to double/triple bonds.

What is a single bond?

Sharing 1 pair of electrons (2 e⁻) between two atoms — represented by one dash (–).

What is a double bond?

Sharing 2 pairs of electrons (4 e⁻) — represented by two dashes (=).

What is a triple bond?

Sharing 3 pairs of electrons (6 e⁻) — represented by three dashes (≡).

What is a bonding pair vs. lone pair?

Bonding pair: shared between two atoms (forms the covalent bond). Lone pair: on one atom only; not shared.

What is a coordinate covalent bond?

A bond where BOTH electrons come from one atom. # coordinate bonds = bonds to central atom − bonding capacity of central atom.

What are the diatomic elements (HOFBrINCl)?

H₂, O₂, F₂, Br₂, I₂, N₂, Cl₂ — plus P₄ and S₈ in elemental form.

What is a polyatomic ion in the context of Lewis structures?

A covalently bonded group of atoms with an overall charge. Add 1 e⁻ per unit of negative charge; subtract 1 e⁻ per unit of positive charge before drawing.

Name a simple binary ionic compound (e.g., CaCl₂)?

Metal name first, then nonmetal with suffix –ide. CaCl₂ = Calcium chloride. Na₃N = Sodium nitride.

Criss-cross rule for ionic formulas?

Write ion symbols with charges; swap the numbers of the charges diagonally to become subscripts; reduce by GCF. (e.g., Al³⁺ + O²⁻ → Al₂O₃)

How do you name a multivalent (transition) metal compound?

Write the metal name, then in parentheses a Roman numeral for its charge, then the anion name. e.g., Cr₂O₃ = Chromium(III) oxide.

How do you find the charge of a multivalent metal from a formula?

Reverse subscripts to get charges; verify anion charge; adjust with multiplier if needed. e.g., Mn(SO₄)₂ → Mn⁴⁺ → Manganese(IV) sulfate.

Name these polyatomic ions: NO₃⁻, SO₄²⁻, PO₄³⁻, OH⁻, NH₄⁺, CO₃²⁻

NO₃⁻ = nitrate; SO₄²⁻ = sulfate; PO₄³⁻ = phosphate; OH⁻ = hydroxide; NH₄⁺ = ammonium; CO₃²⁻ = carbonate

Oxyanion prefix/suffix rules?

+1 O atom → per-…-ate; normal → -ate; –1 O atom → -ite; –2 O atoms → hypo-…-ite. Charge stays the same. (e.g., ClO₄⁻ perchlorate, ClO₃⁻ chlorate, ClO₂⁻ chlorite, ClO⁻hypochlorite)

How do you name a covalent (molecular) compound?

Use Greek prefixes for each element's count (mono- omitted for first element); change suffix of second element to –ide. e.g., CCl₄ = carbon tetrachloride; N₂O₅ = dinitrogen pentoxide.

Greek prefixes 1–10 for covalent compounds?

1=mono, 2=di, 3=tri, 4=tetra, 5=penta, 6=hexa, 7=hepta, 8=octa, 9=nona, 10=deca

How do you name a binary acid?

Prefix hydro- + nonmetal root + -ic acid. e.g., HCl(aq) = hydrochloric acid; HF(aq) = hydrofluoric acid; H₂S(aq) = hydrosulfuric acid.

How do you name an oxyacid (ternary acid)?

Based on the polyatomic anion: –ate anion → –ic acid; –ite anion → –ous acid. e.g., HNO₃ = nitric acid; H₂SO₃ = sulfurous acid; HClO₄ = perchloric acid.

What is a hydrate? How are they named?

An ionic compound with water molecules trapped in the crystal lattice (ionic compound · n H₂O). Named like the ionic compound + Greek prefix + hydrate. e.g., K₃PO₄ · 5H₂O = potassium phosphate pentahydrate.

What is an oxyacid salt? Give an example.

When an oxyacid reacts with a base, the anion may keep extra H⁺. e.g., HCO₃⁻ = hydrogen carbonate; NaHCO₃ = sodium hydrogen carbonate (baking soda).

ΔEN benchmarks for bond type?

0.00–0.39 → nonpolar covalent; 0.40–1.8 → polar covalent; > 1.8 → ionic bond.

How do you calculate ΔEN?

Subtract the smaller EN value from the larger: ΔEN = |EN₁ – EN₂|. Always a positive number.

What is a dipole moment?

A measure of bond polarity — an arrow pointing toward the more electronegative atom (δ–), away from the less electronegative (δ+).

How do lone pairs on the central atom affect polarity?

Lone pairs create a region of concentrated negative charge, making the molecule asymmetrical and contributing to a net dipole — the molecule becomes polar.

Steps to determine if a MOLECULE is polar?

1) Draw Lewis structure. 2) Calculate ΔEN for each bond. 3) Draw dipole arrows. 4) Cancel dipoles that are equal and opposite. 5) Check symmetry. 6) Check for lone pairs on central atom. Net uncancelled dipole = polar molecule.

Is CCl₄ polar or nonpolar? Why?

Nonpolar — even though each C–Cl bond is polar (ΔEN ≈ 0.5), the molecule is symmetrical (tetrahedral), so all dipoles cancel. No lone pairs on C.

Is H₂O polar? Why?

Polar — O has 2 lone pairs that push the H atoms together, making the molecule V-shaped (asymmetrical). Dipoles do NOT cancel; net dipole moment exists.

Classify bond: B & P (ΔEN = 0.15)?

Nonpolar covalent (0.15 < 0.4)

Classify bond: S & O (ΔEN = 1.0)?

Polar covalent (between 0.4 and 1.8)

Classify bond: Li & O (ΔEN = 2.46)?

Ionic (> 1.8)

What are intermolecular forces (IMFs)?

Attractive forces BETWEEN molecules (not within). Much weaker than covalent or ionic bonds (intramolecular forces).

Name the 3 IMFs from weakest to strongest.

1) London Dispersion Forces (LDF) — weakest; 2) Dipole-Dipole Forces; 3) Hydrogen Bonding — strongest.

What are London Dispersion Forces (LDF)?

Temporary dipoles caused by random electron movement that induce dipoles in neighboring atoms/molecules. Present in ALL substances — polar and nonpolar.

What affects the strength of LDF?

1) Size/mass of molecule (more electrons → stronger temporary dipole). 2) Shape (longer molecules have more surface area contact → stronger LDF).

What are dipole-dipole forces?

Electrostatic attraction between the positive end of one polar molecule and the negative end of another. Only in POLAR molecules. Stronger than LDF.

What are hydrogen bonds?

The strongest IMF — a special dipole-dipole force. Requires H bonded directly to N, O, or F (highly electronegative). Strongest when N/O/F–H–N/O/F are in a straight line.

Why are hydrogen bonds so strong?

Large EN difference between H and N/O/F creates a very polarized bond; H is tiny so the positive charge is highly concentrated, allowing close approach.

IMFs and boiling/melting point?

Stronger IMFs → higher boiling point, higher melting point, higher viscosity, higher surface tension, lower vapour pressure, slower evaporation.

Stronger IMFs → higher boiling point, higher melting point, higher viscosity, higher surface tension, lower vapour pressure, slower evaporation.

Polar solutes dissolve in polar solvents (e.g., NaCl in water). Nonpolar solutes dissolve in nonpolar solvents (e.g., wax in hexane). Polar and nonpolar substances do NOT dissolve each other.

Are there IMFs in ionic compounds?

No — ionic compounds are held together by ionic bonds (electrostatic attraction between cations and anions), not IMFs.

Properties of ionic vs. covalent compounds?

Ionic: crystalline solid, high MP, conducts electricity when molten or dissolved. Covalent: solid/liquid/gas, low MP, poor conductor, generally flammable, often insoluble in water.

How do you use IMFs to rank boiling points of substances?

Identify the IMFs present in each molecule. Rank IMFs: LDF < dipole-dipole < H-bonding. Also consider molecular size (larger = stronger LDF). Stronger IMFs = higher boiling point.

NH₄⁺, OH⁻, NO₃⁻, NO₂⁻, CO₃²⁻

ammonium, hydroxide, nitrate, nitrite, carbonate

SO₄²⁻, SO₃²⁻, PO₄³⁻, ClO₃⁻, ClO₄⁻

sulfate, sulfite, phosphate, chlorate, perchlorate

CN⁻, MnO₄⁻, CrO₄²⁻, Cr₂O₇²⁻, C₂H₃O₂⁻

CN⁻, MnO₄⁻, CrO₄²⁻, Cr₂O₇²⁻, C₂H₃O₂⁻

HCO₃⁻, HSO₄⁻, H₂PO₄⁻, HPO₄²⁻

hydrogen carbonate, hydrogen sulfate, dihydrogen phosphate, hydrogen phosphate