SCIENCE 10 FINAL

Exothermic reaction

A reaction that RELEASES energy (heat/light) to surroundings. Gets hot. Energy is a product. Example: combustion, respiration

Endothermic reaction

A reaction that ABSORBS energy from surroundings. Gets cold. Energy is a reactant. Example: photosynthesis, ice packs

Activation energy

The minimum energy required for a reaction to occur. Catalysts lower this.

Energy diagram - exothermic

Reactants start HIGH, products end LOW. Energy is released.

Energy diagram - endothermic

Reactants start LOW, products end HIGH. Energy is absorbed.

Physical change

A change in appearance or state where NO new substance is formed. Usually reversible. Example: ice melting, cutting paper

Chemical change

A change that produces one or more NEW substances. Usually irreversible. Example: burning wood, rusting iron

6 signs of a chemical change

1. Gas produced 2. Heat/light given off 3. Colour change 4. Precipitate formed 5. Odour change 6. Difficult to reverse

Synthesis reaction

A + B → AB. Two or more substances combine to make ONE new substance. Example: 2H₂ + O₂ → 2H₂O

Decomposition reaction

AB → A + B. One substance breaks down into simpler substances. Example: 2H₂O → 2H₂ + O₂

Single displacement reaction

A + BC → AC + B. One element replaces another in a compound. Example: Zn + 2HCl → ZnCl₂ + H₂

Double displacement reaction

AB + CD → AD + CB. Two compounds swap partners. Example: NaCl + AgNO₃ → NaNO₃ + AgCl

Combustion reaction

Fuel + O₂ → CO₂ + H₂O + Energy. A fuel burns in oxygen. Always produces CO₂ and H₂O. Example: CH₄ + 2O₂ → CO₂ + 2H₂O

Acid-base (neutralization) reaction

Acid + Base → Salt + Water. Example: HCl + NaOH → NaCl + H₂O

How to identify synthesis

Multiple reactants, only ONE product

How to identify decomposition

Only ONE reactant, multiple products

How to identify single displacement

One lone element as a reactant

How to identify double displacement

Two compounds react and exchange ions

How to identify combustion

O₂ is always a reactant, CO₂ and H₂O always products

How to identify acid-base

Always involves an acid AND a base together

Reaction rate

The speed at which a chemical reaction occurs — how fast reactants become products

Collision theory

For a reaction to occur, particles must: 1. Collide 2. Have enough energy 3. Collide at the correct angle

Effect of temperature on reaction rate

Higher temperature = particles move faster = more frequent energetic collisions = FASTER rate

Effect of concentration on reaction rate

Higher concentration = more particles in same space = more collisions = FASTER rate

Effect of surface area on reaction rate

Greater surface area = more particles exposed = more contact = FASTER rate

Catalyst

A substance that speeds up a reaction WITHOUT being consumed. Lowers activation energy. Example: enzymes

Inhibitor

A substance that SLOWS DOWN a reaction rate

Effect of pressure on reaction rate

Higher pressure = gas particles pushed closer = more collisions = FASTER rate (gases only)

pH scale

Runs from 0-14. Measures how acidic or basic a solution is. Below 7 = acid, 7 = neutral, above 7 = base

Properties of acids

pH below 7, taste sour, react with metals, turn blue litmus RED, contain H⁺ ions

Properties of bases

pH above 7, taste bitter, feel slippery, turn red litmus BLUE, contain OH⁻ ions

Neutralization

Acid + Base → Salt + Water. pH moves toward 7 (neutral)

Litmus paper in acid

Turns RED

Litmus paper in base

Turns BLUE

Phenolphthalein in acid

Colourless

Phenolphthalein in base

Turns PINK/RED

Universal indicator - red/orange

pH 1-3, strong acid

Universal indicator - green

pH 7, neutral

Universal indicator - blue/purple

pH 11-14, strong base

HCl

Hydrochloric acid. Found in stomach acid. Strong acid.

H₂SO₄

Sulfuric acid. Found in car batteries. Strong acid.

NaOH

Sodium hydroxide. Found in drain cleaner. Strong base.

Organic compound

A compound that contains CARBON (usually bonded to hydrogen). From living things. Generally not water soluble. Often flammable. Example: glucose, methane

Inorganic compound

A compound that generally does NOT contain carbon. From non-living sources. Often water soluble. Example: NaCl, H₂O, Fe₂O₃

Difference between organic and inorganic

Organic = contains carbon, from living things, flammable. Inorganic = no carbon, from minerals, higher melting point

CO₂ - organic or inorganic?

INORGANIC — exception to the rule. Contains carbon but is classified as inorganic.

Lewis dot diagram

A diagram showing the valence (outer shell) electrons of an atom as dots around the element symbol

Valence electrons

The electrons in the OUTERMOST shell of an atom. Equal to the group number for main group elements.

How many valence electrons - Group 1

1 valence electron

How many valence electrons - Group 2

2 valence electrons

How many valence electrons - Group 14

4 valence electrons

How many valence electrons - Group 15

5 valence electrons

How many valence electrons - Group 16

6 valence electrons

How many valence electrons - Group 17

7 valence electrons

How many valence electrons - Group 18

8 valence electrons (full/stable)

Octet rule

Atoms want 8 valence electrons to be stable (except hydrogen which needs 2)

Ionic bond

Formed when a METAL transfers electrons to a NON-METAL. One atom loses electrons (cation +), one gains electrons (anion -)

Covalent bond

Formed when two NON-METALS share electrons. Can be single, double, or triple bonds.

Single covalent bond

Atoms share 1 PAIR of electrons (2 electrons total). Example: H-H

Double covalent bond

Atoms share 2 PAIRS of electrons (4 electrons total). Example: O=O

Triple covalent bond

Atoms share 3 PAIRS of electrons (6 electrons total). Example: N≡N

Ionic vs covalent - elements involved

Ionic = metal + non-metal. Covalent = non-metal + non-metal

Ionic vs covalent - what happens to electrons

Ionic = electrons TRANSFERRED. Covalent = electrons SHARED

Na Lewis diagram

Na has 1 valence electron. In ionic bond, loses it to become Na⁺

Cl Lewis diagram

Cl has 7 valence electrons. In ionic bond, gains 1 to become Cl⁻ (full octet)

NaCl Lewis diagram

Na transfers its 1 electron to Cl. Result: [Na]⁺ [Cl]⁻

H₂O Lewis diagram

O shares one electron with each H. O has 2 lone pairs. Shape: bent/V-shape

Periodic table - period

A HORIZONTAL ROW on the periodic table. Period number = number of electron shells.

Periodic table - group/family

A VERTICAL COLUMN on the periodic table. Group number = number of valence electrons. Same group = similar properties.

Atomic number

Number of PROTONS in the nucleus. Also equals number of electrons in a neutral atom.

Mass number

Number of PROTONS + NEUTRONS in the nucleus

Number of neutrons formula

Mass number MINUS atomic number

Group 1 - Alkali Metals

Very reactive metals. 1 valence electron. React vigorously with water. Example: Li, Na, K

Group 2 - Alkaline Earth Metals

Reactive metals. 2 valence electrons. Example: Mg, Ca

Group 17 - Halogens

Very reactive non-metals. 7 valence electrons. Example: F, Cl, Br

Group 18 - Noble Gases

Unreactive/stable. Full outer shell (8 valence electrons). Example: He, Ne, Ar

Metals - properties

Shiny, good conductors of heat and electricity, malleable, ductile, solid at room temperature (except Hg)

Non-metals - properties

Dull, poor conductors, brittle, lower melting points, many are gases at room temperature

Metalloids

Elements with properties of BOTH metals and non-metals. Found along the staircase line. Example: Si, B, As

Periodic trend - atomic size across a period

DECREASES left to right (more protons pull electrons closer)

Periodic trend - atomic size down a group

INCREASES top to bottom (more electron shells added)

Periodic trend - reactivity of metals

INCREASES going DOWN a group. DECREASES going across a period left to right.

Periodic trend - reactivity of non-metals

INCREASES going UP a group. INCREASES going across a period left to right.

Law of Conservation of Mass

Matter cannot be created or destroyed. Atoms on both sides of an equation must be EQUAL.

Balancing equations - what can you change?

Only change COEFFICIENTS (large numbers in front). NEVER change subscripts.

Coefficient

The large number in FRONT of a formula in an equation. Multiplies ALL atoms in that formula.

Subscript

The small number AFTER an element symbol. Shows how many of that atom are in the compound. NEVER change when balancing.

Steps to balance an equation

1. Write unbalanced equation 2. Count atoms each side 3. Add coefficients starting with most complex molecule 4. Balance metals first, then non-metals, then H, then O last 5. Check all atoms

Balancing - what to do if you get fractions

Multiply ALL coefficients by 2 to clear fractions

Hydroxide ion

OH⁻ Charge: -1

Nitrate ion

NO₃⁻ Charge: -1

Sulfate ion

SO₄²⁻ Charge: -2

Carbonate ion

CO₃²⁻ Charge: -2

Phosphate ion

PO₄³⁻ Charge: -3

Ammonium ion

NH₄⁺ Charge: +1

Balanced equation example - water

2H₂ + O₂ → 2H₂O. Left: 4H 2O. Right: 4H 2O ✓

Balanced equation example - iron oxide

4Fe + 3O₂ → 2Fe₂O₃. Left: 4Fe 6O. Right: 4Fe 6O ✓

Balanced equation example - combustion of propane

C₃H₈ + 5O₂ → 3CO₂ + 4H₂O. Left: 3C 8H 10O. Right: 3C 8H 10O ✓

H symbol and atomic number

Hydrogen. Atomic number 1. Most abundant element in universe.

C symbol and atomic number

Carbon. Atomic number 6. Basis of all organic compounds.