SNC 2D1 Chemistry Notes

WHMIS 2015 and HHPS

• WHMIS stands for Workplace Hazardous Materials Information System.
• It's a system for providing health and safety information on hazardous products in Canadian workplaces.
• WHMIS has 3 parts:
  1. Symbols
  2. Safety Data Sheets (SDSs)
     • SDSs are summary documents about the hazards of a product and safety precautions.
     • They tell users about the hazards, safe usage, what to expect if recommendations aren't followed, how to recognize exposure symptoms, and what to do in emergencies.
  3. Education and Training
     • In Canada, workplaces using hazardous products must have a WHMIS program.
     • Workers must be educated and trained to understand hazards and work safely with hazardous products.

Other Safety Symbols

• Hazardous Household Product Symbols are also used (HHPS).

Safety Questions

1. What often causes accidents in the science classroom?
2. What is the proper way to smell something?
3. Do you ingest anything using lab equipment?
4. What could happen if you light a Bunsen burner incorrectly?
5. What safety equipment must always be worn?
6. How do you know what chemicals you are using?
7. What do you do if you spill an acid or a base on your skin?
8. What do you do if you spill an acid or a base in your eye(s)?
9. Where do you empty contaminated chemical (unless the teacher tells you it’s safe to use the sinks)?
10. What would you do if someone caught on fire?

Lab Equipment

• Beaker
• Tongs
  • Purpose: Pick up hot beakers
• Tweezers
  • Purpose: Pick up small pieces of chemicals
• Test tube rack
  • Purpose: Holds test tubes
• Test tubes
  • Purpose: Hold small amounts of chemicals
• Dropper / Pipette
  • Purpose: Transfer small amounts of liquids
• Mortar & pestle
  • Purpose: Crush substances
• Wash bottle
  • Purpose: Rinse equipment
• Test tube tongs
  • Purpose: Pick up hot test tubes
• Erlenmeyer flask
  • Purpose: Holds or swirls liquids
• Scoopula
  • Purpose: Scoops chemicals
• Florence flask
  • Purpose: Holds liquids for heating
• Graduated cylinder
  • Purpose: Measures liquids accurately
• Rubber stoppers
  • Purpose: Close off test tubes and flasks
• Crucible & cover
  • Purpose: Heat chemicals to high temperatures
• Bunsen burner
  • Purpose: Connects to gas supply for flame
• Watch glass
  • Purpose: Transfers or holds powders
• Evaporating dish
  • Purpose: Heats to evaporate liquids
• Clay triangle
  • Purpose: Holds crucible over flame
• Beaker
  • Purpose: General purpose – holds liquids
• Wire gauze
  • Purpose: Supports beakers over flame
• Burette
  • Purpose: Dispenses liquids very accurately
• Safety goggles
  • Purpose: Protect eyes
• Funnel
  • Purpose: For pouring liquids into smaller containers
• Retort stand
  • Purpose: Supports clamps and other equipment
• Crucible tongs
  • Purpose: Pick up hot crucible
• Glass stirring rod
  • Purpose: Stirs chemicals
• Sparker
  • Purpose: Lights Bunsen burners
• Wire brush
  • Purpose: Cleans glassware
• Test tube clamp
  • Purpose: Holds test tube over flame
• Ring clamp
  • Purpose: Supports other equipment over flame

Physical Property

• A characteristic or description of a substance.
• Does NOT involve forming a new substance.
• e.g. Colour, texture, density, smell, solubility, taste, melting point, physical state, etc.

Chemical Property

• Characteristic behaviour that occurs when the substance changes into something new.
• e.g. reaction of acid with a base, flammability, bleaching ability, corrosion

Physical Change

• Change that does NOT produce a new substance.
• e.g. changes of state, ripping paper, chopping wood
• Many physical changes can be reversed.

Chemical Change

• The change that a substance goes through to produce a new substance
• 5 clues that a chemical change has taken place:
  1. A new colour appears
  2. A solid material (a precipitate) forms in a liquid
  3. Heat or light is produced or absorbed
  4. Bubbles of gas are formed
  5. The change is (generally) difficult to reverse
• The only way to tell for sure that a chemical change has taken place is to conduct further tests on the products
• Some chemical change can be reversed (i.e. rechargeable battery)

Chemical vs Physical Changes Examples

• Rusting iron: Chemical change (new colour, new physical properties)
• Toasting bread: Chemical change (new colour, new odour /gas produced)
• Evaporating water: Physical change (change state)
• Burning a candle: Chemical and Physical change (heat/light , wax melts new odour, change of state)
• Frost on a car window: Physical change (change of state)
• Turning a light bulb on: Physical change (light produced)
• Boiling water: Physical change (change in state)
• Formation of clouds: Physical change
• A firefly glowing: Chemical change (light produced)
• Breaking a stick: Physical change
• Wax melting: Physical change
• Frying an Egg: Chemical change (gas produced/new odour)

Teacher Demos

• phenolphthalein + Sodium hydroxide: new colour - Chemical change
• iron (III) chloride + Potassium thiocyanate: new liquid colour - Chemical change
• ammonium nitrate + water: dissolving - heat is absorbed
• magnesium + oxygen: bright white light - Chemical change
• Acetic acid + sodium hydrogen carobonate: bubbles - Chemical change
  Sodium iodide + Lead (II) nitrate - solid

Physical vs. Chemical Properties

• Physical: blue colour, density, solubility, melting point
• Chemical: flammability, reactivity with acid/water/base

Physical vs. Chemical Changes

1. Sodium hydroxide dissolves in water: Physical
2. Hydrochloric acid reacts with potassium hydroxide to produce a salt, water, and heat: Chemical
3. A pellet of sodium is slice in two: Physical
4. Water is heated and changed to steam: Physical
5. Potassium chlorate decomposes to potassium chloride and oxygen gas: Chemical
6. Iron rusts: Chemical
7. When placed in water, a sodium pellet catches on fire as hydrogen gas is released and sodium hydroxide forms: Chemical
8. Evaporation: Physical
9. Ice melting: Physical
10. Milk sours: Chemical
11. Sugar dissolves in water: Physical
12. Wood rotting: Chemical
13. Pancakes cooking on a griddle: Chemical

Patterns and the Periodic Table

• Element: a pure substance that cannot be broken down into simpler substances
• Period: a row of elements in the periodic
• Group (Family): a column of elements in the periodic table with similar properties
  • Alkali metal (Group 1): soft, highly reactive metal
  • Alkaline earth metals (Group 2): light, reactive
  • Halogens (Group 17): one of the most reactive group
  • Noble gas (Group 18): very stable, rarely react with any other chemical

Atomic Structure

• Proton:
  • Electrical charge: Positive
  • Symbol: $p+$
  • Location: Nucleus
• Neutron:
  • Electrical charge: Neutral
  • Symbol: $n0$
  • Location: Nucleus
• Electron:
  • Electrical charge: Negative
  • Symbol: $e-$
  • Location: Orbit around the nucleus
• The elements of the periodic table are arranged in order of increasing atomic number (= number of proton)
• Atoms are electrically neutral, where number of proton equals to number of electrons

Standard Atomic Notation

• Write the chemical symbol of the atom and place the atomic number to the lower left and the mass number to the upper left

Bohr-Rutherford Diagram:

• # of Protons = atomic number
• # of Electrons = atomic number
• # of neutron = atomic mass – atomic number

Electron Arrangement and Reactivity

• Valence Electron: electrons that are on the outermost ring
  • Responsible for the element’s reactivity
  • Noble gases are stable because they have completely filled their outer orbits
• Compound: substances made up of two or more elements in a fixed ratio
  • Compounds are formed when elements have their outermost orbit filled

First 20 Elements

Electron Arrangement and Reactivity (cont.)

• lose 1 electron = +1 charge
• lose 2 electrons = +2 charge
• lose 3 electrons = +3 charge
• gain 3 electrons = -3 charge
• gain 2 electrons = -2 charge
• gain 1 electron = -1 charge
• The period number = # of Orbitals
• The column number = # of valence electrons

Lewis Dot Diagram, Atoms and Ions

• Valence electrons are the electrons that are at the outermost ring
• Write the chemical symbol
• Put the number of valence electrons around the symbol
• Isotope: Is an atom with the same number of protons and a different number of neutrons
• Ion: a charged particle that results when an atom gains or loses one or more electrons

Stable Octet Rule

• atoms gain or lose electron(s) to achieve a full outer orbit (usually 8), similar to the nearest noble gases.
• Sodium atom will lose the one electron on the outermost ring. The sodium ion now has the same stable electron arrangement as a neon atom
• Fluorine tends to gain one electron to achieve stable octet

How and Atom Becomes and Ion

• Aluminum lose 3 electrons
• Chlorine gain 1 electrons
• Beryllium lose 2 electrons
• Oxygen gain 2 electrons

Hydrogen

• Can achieve stability by gaining one electron to fill its only orbit (H-)
• Or loses its only electron to form an ion with an ionic charge of +1

Cation

• A positively charged ion
• Metals form cations
• Naming of positive ion is the same
• e.g. $Mg^{2+}$ is still magnesium

Anion

• A negatively charged ion
• Non-metals form anions
• Naming of the negative ion is determined by adding “ide” to the stem of the name
• Hydrogen -> hydride
• Oxygen -> oxide
• Sulfur -> sulfide
• Fluorine -> fluoride
• Chlorine -> chloride
• Nitrogen -> nitride

Ions Practice

Ionic Bonding Using Lewis Structures

For each of the following:
Draw the ionic bond formed using Lewis Structures,
Name the cation,
Name the anion,
State the chemical formula for the ionic compound which is formed,
Name the ionic compound.

• Lithium and Fluorine
  • lithium fluoride (LiF)
• Calcium and Oxygen
  • calcium oxide (CaO)
• Magnesium and Bromine
  • magnesium bromide ($MgBr_2$)
• Sodium and Sulfur
  • sodium sulfide ($Na_2S$)
• Calcium and Nitrogen
  • Calcium nitride ($Ca<em>3N</em>2$)

Ionic Compound

• Metals and non-metals combine to form Ionic compounds by transferring electrons
• Metals: lose electrons to form cations (positive ion)
• Non-metals: gain electrons to form anions (negative ion)
  • Since opposite charges attract, metal and non-metal ions now attract each other
  • the resulting ions all have stable, filled outer electron arrangements as the nearest noble gas

Ionic Compound:

• compounds that are made up of positive and negative ions

Ionic bond:

• the simultaneous strong attraction of positive and negative ions in an ionic compound

Properties of Ionic compound

• hard, brittle solids with high melting points
• Most ionic compounds are electrolytes – dissolve in water to produce a solution that conducts electricity
  • Pure water is a poor conductor of electricity but tap water, lake water, and seawater are good conductors because they contain ions from a variety of sources

Naming Ionic Compounds:

• The first part refers to the metal ion in the compound and the second to the non-metal ion
• Metal ion: name remains the same
• Non-metal ion: ending changes to “-ide” e.g. magnesium chloride

Chemical Formulas:

• In a compound, the number of positive charges must equal the total number of negative charges
• e.g. Magnesium and Chlorine

1. Write the symbol of the elements. Metal on the left, non-metal on the right
2. Add ionic charge of each ion above the symbol
3. Crisscross Rule: Crisscross the numbers of the ionic charges so that they now become subscripts for the opposite elements
4. State the chemical formula and chemical name e.g. Aluminum and oxygen

• Aluminum and oxygen
  • $Al<em>2O</em>3$ (aluminum oxide)
• Magnesium and Sulfur
  • $MgS$ (magnesium sulfide)

Element with Multiple Ionic Charges

• Some metals have two stable cations
• The Roman numerals in rounded brackets are used to indicate the ionic charge of the metal

• Copper (II) nitride
  • $Cu<em>3N</em>2$

Determine the chemical name of PbO2

1. Reverse crisscross
2. Check the charge of the anion. If it has reduced, multiply the cation and anion’s charges by the reduced factor

• Copper (1) bromide

Manganese - (IV) oxide

Multivalent Ions Practice Questions

Use the reverse cross-over method to find the charge on the cation in these compounds

• $Cu_2S$
• $Fe<em>2O</em>3$
• $PbO_2$
• $NiCl_2$
• $CrN$
• $HgO$

Complete the following chart by naming the ionic compounds

• $FeCl_3$
• FeO
• $Cu_2S$
• $PbO_2$

Write the chemical formula for each of the following compounds

• Copper (I) oxide
• Lead (IV) bromide
• Iron (III) sulfide
• Nickel (III) fluoride
• Manganese (IV) sulfide

Polyatomic Ions

• Polyatomic ions are groups of atoms that tend to stay together and carry an overall ionic charge.
  • e.g.The nitrate ion
  • The sulfate ion
• When a compound containing a polyatomic ion is dissolved in water, the metal ion separates from the polyatomic ion, but the atoms of the polyatomic ion stay together as a unit.

Common Polyatomic Ions and Their Ionic Charges

• The suffix –ate in a polyatomic ion denotes the most common number of oxygen atoms

WRITING FORMULAS FOR POLYATOMIC COMPOUNDS

• Since polyatomic ions have an overall charge and these molecules tend to stick together, they can bond with metals to form ionic compounds. To name them, we use the same methods as we use with naming other ionic compounds.

• What is the formula for the ionic compound formed by sodium and a sulfate ion?

  1. Write the symbols of the metal and of the polyatomic group.
  2. Write the ionic charges above the symbols and crisscross them.
  3. Write the formula using subscripts.
     What is the formula for the ionic compound formed by magnesium and a phosphate ion?

Practice: Write the chemical formula for the following:

• Beryllium sulfate
• Aluminum hydroxide
• Lead (II) nitrate
• Copper (II) phosphate

NAMING POLYATOMIC COMPOUNDS

• The name is simply a combination of the name of the metal and the name of the polyatomic ion.
• Name the following ionic compounds.
  • $KClO_3$
  • $Na<em>2SO</em>4$
  • $CaCO_3$
  • $CuNO_3$

OXYACIDS

• Oxyacids are compounds formed when hydrogen combines with polyatomic ions that contain oxygen.
• The hydrogen has an ionic charge of 1+ in these compounds.

Common Oxyacids

Ion name | Ion formula | Ionic charge | Oxyacid formula | Oxyacid name
nitrate | $NO<em>3$ – | 1– | $HNO</em>3$ | nitric acid
chlorate | $ClO<em>3$ – | 1– | $HClO</em>3$ | chloric acid
carbonate | $CO<em>3^{2}$ – | 2– | $H</em>2CO<em>3$ | carbonic acid sulfate | $SO</em>4^{2}$ – | 2– | $H<em>2SO</em>4$ | sulfuric acid
phosphate | $PO<em>4^{3}$ – | 3– | $H</em>3PO_4$ | phosphoric acid

NAMING OXYACIDS

• Name an acid with a polyatomic ion ending with the –ate suffix, drop the –ate suffix and add the suffix, –ic acid.
• Naming Ionic Compounds Worksheet
  • $CaCO_3$
  • KCl
  • $FeSO_4$
  • LiBr
  • $MgCl_2$
  • $FeCl_3$
  • $Zn<em>3(PO</em>4)_2$
  • $NH<em>4NO</em>3$
  • $Al(OH)_3$
  • $PbSO_4$
  • $NaClO_3$
  • $Fe<em>2O</em>3$
  • $(NH<em>4)</em>3PO_4$
  • $HgCl_2$

Formulas of Polyatomic Compounds: Criss-Cross Method

Write the formulas of the compounds produced from the list of ions.

$ClO<em>3$ - $CO</em>3^{2-}$

$OH-$

$SO_4^{2-}$

$PO_4^{3-}$

$NO_3$ -

N$a^+$

$NH_4 +$

$K^+$

$Ca^{2+}$

Lesson 8: Molecules and Covalent Bonding

Main ideas:

1. Bonding between non-metals
2. Sharing of electrons

• When two non-metals bond with each other, both nuclei form strong attractions for the other’s electrons
• Neither atom attracts the other’s electrons strongly enough to pull them away completely
• “Tug of war” of electrons that neither atom ever wins two atoms share each other’s electrons, a bond that holds the atoms together

Covalent bond

• a bond that results from the sharing of outer electrons between non-metal atoms

Molecule

• a particle in which atoms are joined by covalent bonds

Diatomic molecule:

• a molecule consisting of only two atoms of either the same or different elements

Common elements that exist as diatomic molecules.

• Memory Aid: HOFBrINCl the clown
• e.g. $PCl_3$
  1. Write the names for both elements in the same order. Replace the ending of the second element with “-ide”
  2. Add Prefixes.

Covalent Compounds Practice: