Chemistry 10: Study Guide Flashcards

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WHMIS

Workplace Hazardous Materials Information System

Purpose of WHMIS

Protect workers by giving standardized information on hazards

SDS or MSDS (Safety Data Sheet)

Detailed document for each hazardous product giving composition

Flame pictogram

Indicates flammable gases

Flame over circle pictogram

Indicates oxidizing substances that can cause or intensify fire.

Exploding bomb pictogram

Indicates explosives

Gas cylinder pictogram

Indicates compressed

Corrosion pictogram

Indicates corrosive to metals and causes severe skin burns and eye damage.

Skull and crossbones pictogram

Indicates acute toxicity that can be fatal or toxic even in small amounts.

Health hazard pictogram

Indicates long-term health effects like carcinogen

Exclamation mark pictogram

Indicates irritant

Environment pictogram

Indicates aquatic environmental toxicity (not mandatory in WHMIS 2015).

Classroom safety: eye and hair

Always wear goggles and tie back long hair during lab activities.

Classroom safety: food and drink

Do not eat or drink in the lab and never taste chemicals.

Classroom safety: instructions

Read all instructions before starting and follow teacher directions exactly.

Classroom safety: spills and injuries

Report all spills

Classroom safety: disposal

Dispose of chemicals only as instructed

Household hazard shapes

Octagon means danger

Household hazard pictures

Explosive

Physical change

Change in state

Examples of physical change

Ice melting

Chemical change

Change where new substances with different properties form because atoms rearrange.

Examples of chemical change

Iron rusting

Indicator: colour change

Colour changes in a way that is not just dilution

Indicator: gas produced

Formation of bubbles

Indicator: precipitate

Formation of a solid in a solution that was previously clear.

Indicator: energy change

Temperature increase or decrease

Exothermic reaction

A reaction that releases energy to the surroundings and usually feels warm or hot.

Exothermic energy placement

Energy is on the product side of the chemical equation and surroundings’ temperature increases.

Exothermic examples

Combustion of fuels

Endothermic reaction

A reaction that absorbs energy from the surroundings and often feels cold.

Endothermic energy placement

Energy is on the reactant side of the chemical equation and surroundings’ temperature decreases.

Endothermic examples

Photosynthesis

Reaction rate: temperature

Higher temperature makes particles move faster

Reaction rate: concentration

Higher concentration means more particles in the same volume and more collisions.

Reaction rate: surface area

Smaller pieces or powders have more exposed surface area

Reaction rate: catalyst

A substance that speeds up a reaction by lowering activation energy without being used up.

Atomic number

Number of protons in an atom

Period on periodic table

A horizontal row where elements have the same number of occupied electron shells.

Group or family

A vertical column where elements have similar chemical properties and the same number of valence electrons.

Group 1: alkali metals

Very reactive metals with 1 valence electron and a common charge of +1.

Group 2: alkaline earth metals

Reactive metals with 2 valence electrons and a common charge of +2.

Transition metals

Groups 3–12

Group 17: halogens

Very reactive nonmetals with 7 valence electrons and a common charge of −1.

Group 18: noble gases

Very unreactive gases with full valence shells

Metals

Elements that are usually shiny

Nonmetals

Elements that are often dull

Metalloids

Elements along the staircase that have properties between metals and nonmetals (such as silicon and germanium).

Common charge: Group 1

Form +1 ions like Na⁺ and K⁺.

Common charge: Group 2

Form +2 ions like Ca²⁺ and Mg²⁺.

Common charge: Group 13

Form +3 ions like Al³⁺.

Common charge: Group 16

Form −2 ions like O²⁻ and S²⁻.

Common charge: Group 17

Form −1 ions like Cl⁻

Diatomic elements list

Elements that exist as X₂ when pure: H₂

Protons in an atom

Number of protons equals the atomic number of the element.

Electrons in a neutral atom

In a neutral atom

Neutrons in an atom

Neutrons equal the mass number minus the atomic number.

Example: carbon-12

Atomic number 6 → 6 protons and 6 electrons

mass number 12 → 6 neutrons.

Example: sodium-23

Atomic number 11 → 11 protons and 11 electrons

mass number 23 → 12 neutrons.

Steps for molecular (molar) mass

Count atoms of each element

Example: water molar mass

For H₂O

Ionic compound definition

Compound formed by electron transfer between metal and nonmetal (or polyatomic ions) creating cations and anions.

Ionic compound properties

High melting points

Molecular compound definition

Compound formed when nonmetals share electrons

Molecular compound properties

Oftentimes lower melting points and many are gases or liquids at room temperature.

Recognizing ionic from formula

If the formula starts with a metal or NH₄⁺

Recognizing molecular from name

If the name uses Greek prefixes like mono-

Naming simple ionic compounds

Name the metal first and change the nonmetal ending to “-ide” (for example

Naming ionic with polyatomic ions

Keep the polyatomic ion name such as sulfate

Transition metal naming

Determine the metal’s charge and show it with a Roman numeral in brackets (for example

Example FeCl₂ name

Two Cl⁻ (−2 total) means Fe must be +2

Example FeCl₃ name

Three Cl⁻ (−3 total) means Fe must be +3

Naming molecular compounds

Use Greek prefixes to show the number of atoms and end the second element with “-ide.”

Common prefixes list

Mono-

Molecular name example: CO

Carbon monoxide (not “monocarbon monoxide”).

Molecular name example: CO₂

Carbon dioxide.

Molecular name example: N₂O₄

Dinitrogen tetroxide.

Writing ionic formulas

Write charges for cation and anion

Example: magnesium chloride formula

Mg²⁺ and Cl⁻ combine as MgCl₂ because two Cl⁻ balance +2.

Example: aluminum oxide formula

Al³⁺ and O²⁻ form Al₂O₃ so that total charge is zero.

Polyatomic ion in formulas

Use brackets when more than one polyatomic ion is needed

Writing molecular formulas from names

Use the prefixes as subscripts

Example: sulfur trioxide formula

Sulfur trioxide is SO₃.

Example: carbon tetrachloride formula

Carbon tetrachloride is CCl₄.

Law of Conservation of Mass

In a chemical reaction matter is not created or destroyed and total mass of reactants equals total mass of products.

Balancing equations rule

Only change coefficients

Coefficient meaning in equations

A coefficient multiplies all atoms in the formula

General balancing strategy

Balance metals first

Balancing example Fe + O₂

Balanced equation is 4Fe + 3O₂ → 2Fe₂O₃ with equal Fe and O atoms on both sides.

Acid properties

Acids taste sour

Acid formula patterns

Acid formulas usually start with H or end with COOH in many organic acids.

Base properties

Bases taste bitter

Base formula patterns

Bases often contain OH⁻ such as NaOH or KOH or are metal oxides and hydroxides.

pH scale description

Scale from 0 to 14 measuring how acidic or basic a solution is.

Neutral pH value

Neutral solutions like pure water have a pH of 7.

Acidic pH range

Acidic solutions have pH less than 7

Basic pH range

Basic solutions have pH greater than 7

Strength ranges on pH scale

pH 0–3 strong acids