Science 10

Page 1: Science 10: Unit One

  • Chemistry Overview

Page 2: Safety Hazard Symbols

  • Hazard Symbols: indicate both the type of hazard and the degree of hazard.

  • Types of symbols include:

    • Flammable

    • Toxic

    • Explosive

    • Corrosive

Page 3: WHMIS

  • Workplace Hazardous Materials Information System (WHMIS) provides symbols for hazardous materials:

    • Compressed gas

    • Dangerously reactive material

    • Oxidizing material

    • Poisonous and infectious material causing immediate and serious toxic effects

    • Flammable and combustible material

    • Biohazardous infectious material

    • Corrosive material

    • Poisonous and infectious material causing other toxic effects

Page 4: Properties & Classification of Matter

  • Matter: anything that takes up space and has mass.

  • States of Matter: 3 basic states

    • Solid: definite shape and volume

    • Liquid: definite volume, no definite shape

    • Gas: no definite shape or volume

  • Physical Property: can be observed without changing the matter (e.g., size, shape, color, texture).

    • Includes: melting/freezing point, boiling point, solubility, conductivity.

  • Chemical Property: how matter reacts to form new substances (e.g., flammability, ability to rust).

Page 5: Physical vs Chemical Change

  • Physical Changes: change physical properties without changing chemical identity.

  • Chemical Changes: alter the chemical identity of matter (chemical reactions). Includes:

    • Production of a new substance

    • Color change

    • Odor change

    • Change of state

    • Change in energy/temperature

Page 6: Subcategories of Matter

  • Classification

    • Pure Substances: identical particles; constant properties (elements and compounds).

    • Elements: can’t be broken down.

    • Compounds: can be broken down; made of 2 or more elements.

    • Mixtures: combination of two or more pure substances.

      • Heterogeneous Mixture: different composition, substances visible.

      • Homogeneous Mixture: uniform composition, substances not visible.

Page 7: Atoms & Atom Discovery

  • Particle Model of Matter: everything is made up of atoms.

  • Subatomic Particles:

    • Electrons: negatively charged, around nucleus.

    • Protons: positively charged, in nucleus.

    • Neutrons: neutral, in nucleus.

  • Key contributors to atomic theory: Dalton, Thomson, Rutherford, Bohr.

Page 8: Scientist Contributions

  • Dalton: all matter is made of atoms.

  • Thomson: discovered the electron.

  • Rutherford: proposed atomic structure (nucleus with protons and neutrons).

  • Bohr: electrons orbit the nucleus.

Page 9: Atomic Energy Levels

  • Electrons occupy specific energy levels, regions near the nucleus.

  • Ground state: lowest energy (nearest the nucleus).

  • Excited state: higher energy (further from nucleus).

  • Capacity of energy levels: first can hold 2, next two levels can hold 8.

Page 10: Atomic Number, Mass Number & Atomic Molar Mass

  • Atomic Number: number of protons in an element.

  • Neutral atoms have equal protons and electrons.

  • Isotopes: atoms of the same element with different neutron numbers. Mass number = protons + neutrons.

  • Atomic Molar Mass: average mass of an element’s isotopes.

Page 11: The Periodic Table

  • Dmitri Mendeleev organized elements by increasing atomic number and similar properties.

  • Rows: called periods.

  • Columns: called groups or families.

  • Categories: metals, non-metals, and metalloids.

Page 12: Families on the Periodic Table

  • Group 1: Alkali metals (highly reactive).

  • Group 2: Alkaline earth metals (reacts with oxygen).

  • Group 7: Halogens (very reactive non-metals).

  • Group 8: Noble gases (unreactive).

Page 13: Ions

  • Ionization: process of gaining or losing electrons.

  • Ions: charged atoms/groups of atoms.

  • Types of Ions:

    • Cations: positively charged, usually metals.

    • Anions: negatively charged, usually non-metals. Named by changing ending to "ide."

Page 14: Stability and Electron Configuration

  • Atoms aim for stable electron configuration like noble gases.

  • Valence Electrons: determine reactivity; same number in a family/group.

Page 15: The Octet Rule

  • Atoms aim for a full outer energy level (8 electrons).

  • Exceptions: H, Li, Be need 2 electrons (config like He).

Page 16: Ionic Compounds

  • Formed through electron transfer between metals and non-metals.

Page 17: Properties of Ionic Compounds

  • Ionic compounds are:

    • Solid at room temperature

    • Soluble in water

    • Conduct electricity

Page 18: Naming & Writing Formulas for Ionic Compounds

  • Naming convention: cation name first, anion name second (change ending to "ide").

  • Formula: reflects ratio of ions for neutral charge (drop & swap method).

Page 19: Drop & Swap Method

  • Example: Mg2+ + CI- → MgCl2 (calculation of ion charges).

Page 20-23: [Empty Pages]

Page 24: Molecular Compounds

  • Formed by sharing electrons between non-metals to form molecules.

  • Tend to be poor conductors of electricity.

Page 25: Molecular Elements

  • Diatomic Elements: molecules with two identical atoms (e.g., N2).

  • Naming: Greek prefixes indicate atom quantity; second element ends with "ide."

Page 26: Prefixes for Naming Binary Molecular Compounds

  • Mono-: 1

  • Di-: 2

  • Tri-: 3

  • Tetra-: 4

  • Penta-: 5

  • Hexa-: 6

  • Hepta-: 7

  • Octa-: 8

  • Ennea-: 9

  • Deca-: 10

Page 27-28: [Empty Pages]

Page 29: Solubility

  • Solubility: varies between ionic compounds (some dissolve better).

  • Precipitate: solid formed from a solution with low solubility.

Page 30: Acids & Bases

  • Acids: pH < 7, produces H+ ions in water.

  • Bases: pH > 7, produces OH- ions in water.

  • Examples: vinegar (acid), soap (base).

  • Buffers maintain pH levels.

Page 31: Acid-Base Tests

  • Methods for testing:

    • Litmus paper (red for acid, blue for base).

    • pH measurement.

    • Indicators change color in presence of acids/bases.

Page 32: Neutralization

  • Neutralization: reaction between acids and bases produces water and salt.

  • Common reactions (bee stings, heartburn).

Page 33: Chemical Reactions

  • Chemical Change: substances react to form new substances (reactants → products).

  • Energy flow: exothermic (releases energy) or endothermic (absorbs energy).

Page 34: Characteristics of a Chemical Reaction

  • New substances produced, change in state, color/odor change, energy flow.

  • Consistent with the Law of Conservation of Mass.

Page 35: Writing Chemical Equations

  • Chemical equations represent reactions in symbolic form.

  • Word equations: describe reactions (e.g., Magnesium + Hydrochloric acid → Magnesium chloride + Hydrogen gas).

Page 36: Balancing Equations

  • Balanced equations: equal number of atoms on both sides.

  • Coefficients adjust amounts; subscripts should not change.

Page 37-40: Balancing Strategies

  • Strategies for balancing equations include:

    • Balancing noticeable differences first, using common multiples, treating polyatomic ions as groups.

Page 41: Types of Chemical Reactions

  • Five Common Types:

    • Synthesis

    • Decomposition

    • Single Replacement

    • Double Replacement

    • Hydrocarbon Combustion

Page 42: Strategies for Completing Equations

  • Identify reaction type (combustion, decomposition, etc.), use drop & swap strategy for ionic compounds, remember to balance.

Page 43-45: Additional Strategies for Completing Equations

  • Various examples for applying strategies for different reaction types.

Page 46: The Mole (mol)

  • Mole: SI unit for the amount of substance. 1 mol = 6.02 x 10^23 particles.

Page 47: Calculating Molar Mass

  • Molar mass: atomic mass from the periodic table. Add masses of each element in a compound for total molar mass.

Page 48-50: [Empty Pages]

Page 51: Atoms & Moles

  • Relation between atoms and moles through Avogadro’s number.

  • Key equations for calculating moles and mass.