# SCH3U Grade 11 Chemistry Nomenclature, Trends, Reactions Test

Grade 11 Chemistry 11 U Notes – Trends of the periodic Table

Significant Digits:

1. All non-zero digits are significant.
2. Leading and following zeros used to indicate the position of the decimal are not significant
3. All zeros between two other digits are significant.
4. All zeros that follow a non-zero AND is to the right of the decimal point are significant

Scientific notation:

-If the number is greater than one,

 Move the decimal point to the left until there is just one digit in front of it.

 Count the number of decimal places to be written positively

-If the number is less than one,

 Move the decimal to the right until there is just one number in front of it (cannot be zero).

 Count the number of places the decimal place has been moved and that is the negative exponent or power of 10 to be written

-When multiplying, you are to multiply decimals and exponents separately.

-When adding/subtracting, convert each exponent to the same exponent

-When moving decimal place left, the exponent adds one

-When moving decimal place left, the exponent subtracts one

SI Units

Meters, Seconds, Kilogram, Mole

Pico: 1 x 10^-12

Centi: 1 x 10^-2

Conversion of units.

The Atom:

Cloud Model: electrons take up space around the nucleus in a cloud.

-Max number of electrons per shell:

1st: 2 electrons

2nd: 8 electrons

-Math equation for finding max electrons per shell

2n^2 when n represents the orbital #

Nucleus

Proton – Positive

Neutrons – Neutral

-Same mass as each other

Carbon example:

6 Protons = atomic number

6 electrons = atomic number = proton

6 neutrons = atomic mass – atomic number

Ion: gain or lose electrons to have full outer orbital

Carbon 12 -> Atomic Mass

Carbon 13 -> gains one electron, becomes an isotope

-Becomes C

Radioisotopes: isotopes that are unstable and are radioactive, meaning they decay they release energy and subatomic particles

-Occurs if nucleus is unstable.

Average Atomic Mass: used to determine the average mass found on periodic table

Isotope Abundance (%): how often a particular isotope occurs naturally

Question: Copper 63: 69% and mass of 69.9 u, Copper 65 occurs 30.8% and mass of 64.93 u.

Average Atomic Mass copper = (atomic mass of Copper 63 x % of abundance copper 63) + (aM of copper 65 x % abundance of copper 65)

Trends in the periodic Table

Atomic Radius: distance from nucleus to the region where electrons live.

Law of Attraction

-Opposites attract

-Like charges repel

-number of protons increase, electrons will increase but be attracted to a confined region of the shell.

Shielding Effect: the effect of shielding valence electrons with filled shells

Effective Nuclear Charge: the apparent nuclear charge as experienced by the outer most electron shell, low as result of inner shielding (spaced farther apart)

Ionization energy: energy required to remove one electron from an atom or ion.

A(g) + energy -> A+ + electron

-With each removal of electrons, the next will become even stronger because the force of attraction will be stronger for each lesser electron

-As you move across, IE will increase, as they’re more stable

-As you move down, IE will decrease as distances increase and shielding takes place

Electron Affinity: energy released or absorbed when electrons are added

If Energy is released when electrons are added, EA is a negative integer and creates an anion.

If Energy is added when electrons are added, EA is a positive number and creates a cation.

Electronegativity: an indicator of ability to attract electrons

-shared electrons for Li and Br for example, will share, but Br will be closer to core than Li because of it’s atomic radius.

Equal Sharing Compound – diatomic molecules

-with 2 identical atoms, equal sharing is accomplished.

Chemical Bonding

“bonding involving the interaction of valence electrons between atoms which usually will create a compound that is more stable”

• Lewis-Dot Structure:diagram of valence electrons only
• Atoms are labelled clockwise, one at a time until each pole has 2
• All elements want to become stable
• GAINS, LOSE, or SHARE electrons to accomplish that
• Always wants to have a full valence shell

Octet Rule: allows us to predict what bonds fill form between atoms when they bond.

• Valence electrons will bond in such a way that will fill up all 8 valence electrons

Isoelectric: happens when 2 atoms have the same electron configuration

• Eg: Na+ is isoelectric to Ne (the ion is identical as a normal Ne valence shell)
• Eg: Cl is isoelectric to Argon

Ionic Bonds

• Forms when there is an attraction between oppositely charged ions
• Formation of an ionic bond must have net charge of zero
• Generally between metals and non-metals
• Determining the type of bond using difference (delta) in electronegativity (EN)
• Higher EN – Lower EN (more than one of an atom doesn’t matter)
• If Delta EN is 1.7+, it’s an ionic bond.
• If Delta EN is 0-0.5, it’s a pure covalent bond
• If Delta EN is 0.5 – 1.7, it’s a polar covalent bond
• Eg: Na + Cl
• The one Na valence goes to Cl to complete Cl’s octet rule
• [Na]+ [Cl] -> Na+1 + Cl-1 = 0
• Transition Metals(with multiple ion possibilities)
• 2 Irons + 3 Oxygens
• Fe2+ or Fe3+ + 3(O-2)  (O with filled electrons)
• 2(3+) + (-6) = 0
• Properties of ionic bonds
• Much stronger attraction due to +ive and –ive bonds
• Each ion is surrounded by an ion of the opposite charge in a solid model
• Because each charge is extremely strong, pulling other together
• In solutions, they’re dissolved, and split apart

Covalent Bonds

• These types create bonds by sharingelectrons
• Each atom contributes one or more electron to the compound
• Eg: H2 shares it’s valence electron in the middle to fulfill the octet rule for each
• Pure Covalent Bonds: when electrons are equally shared
• Diatomic molecules are all pure covalent with delta EN of zero
• Eg: CCl4 -> 3 – 2.5 = 0.5  -> Covalent bond
• Eg: Cl2O -> 3.5 – 3 = 0.5  -> Covalent bond
• Polar Covalent Bonds: when there’s an unequal sharing of electrons
• one atom will attract more electrons than another
• Eg: CF4 Delta EN = 4.0-2.5 = 1.5  -> polar covalent
• If one EN value of an atom is higher than another, it means that it will pull more electrons to that side*
• If more attraction is to that atom, it’s partially negative to that side (d-) (delta)
• If more attraction is less attracted to that atom, it’s partially positive (d+) (delta)
• Intramolecular Forces: attraction between atoms
• Intermolecular Forces: attraction between molecules
• Have very strong attraction, but no where near as high as ionic
• Have higher boiling and melting points than non-polar molecules
• Multiple Covalent Bonds
• when more than 1 electron is being shared
• Single Bond: if one electron is shared until octet rule is satisfied
• Double Bond: must look at 2 atoms at the same time, when 2 electrons are shared
• Triple Bond: when 3 electrons are shared
• Non – Polar compounds
• These compounds do not have positive or negative ends because EN of surrounding molecules cancels them out
• usually true for symmetrical molecules
• Comparing Ionic and Molecular Compounds
 Properties Ionic Molecular Melting Point High Low Conducts in solid No No Conducts in solution Yes No Solubility Yes in water Polar in polar solvents, non polar in non polar solvents State Solid Any Consistency of solid Brittle, hard Flexible, soft

Binary Nomenclature

• IUPAC Ionic Nomenclature
• all lower case
• metals first
• keeps root name, then add ide prefix at last word
• eg: sodium phosphide = Na2P
• Find charges by doing criss cross, then simplify the bond

• Polyatomic Ions
• For naming pre-made polyatomic ions
• Combine criss cross and do charges
• Use memorized names
• Use charges given to it
• Binary Molecular Compounds
• Use greek prefixes
• mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca
• 1) higher group goes first
• 2) If same group, higher period goes first
• 3) First element name doesn’t change, just add prefix
• 4) Second element add suffix –ide
• Hydrogen binary acids
• In IUPAC: add “aqueous” or “(aq)”
• Use “Acid” ending in Classical Name
• The compound will NOT have oxygens in it
• 1) omit hydrogen, and put in hydro
• 2) add suffix “ic” to the end
• Variations in oxyacids
• different oxygens based on endings
• using –ATE oxygen polyatomic ion as reference, we have:
• Per- : add 1 oxygen
• -ite  : take one oxygen
• hypo- ___ -ite : take 2 oxygen
• Group 16 and 15 only ate and ite endings!
• Group 14 only ate endings!
• Oxyacid naming
• –ite ending polyatomic ion -> -ous acid
• –ate ending polyatomic ion -> -ic acid

Binary Nomenclature

• Metals with non-constant valence charges
• Roman Numeral Nomenclature: “stock System”/IUPAC
• use roman numerals to denote their charges used
• Iron (II) oxide
• Greek Latin names (Classical System)
• Use of greek names and latin endings
• need memorization
• Balancing Equations
• The goal of balancing equations is to have them the same mass on both sides
• Law of conservational mass: both reactant and product must have same number of atoms
• Use coefficients ONLY to balance equations
• Hydrogenated Polyatomic Ions
• combining polyatomic ions with Hhydrogen ions
• Name them by having polyatomic ion, then adding a hydrogen
• Eg K2HPO3 -> potassium hydrogen phosphite
• The hydrogen uses the greek prefix system
• Hydrogen is not affected by bracketed amounts
• When converting to formula, the charges will matter again
• Charges will change, depending on number of that ion
• Must manually combine hydrogen and polyatomic molecule with correct charges to result to overall 1 or -1
• Hydrates
• uses greek letters to represent how many water (Hydrates) are needed
• In chemical formula, water molecule is separated by a dot

Types of Chemical Reactions

• Synthesis
• element to element
• Metal to non-metal
• simple ionic reaction
• Univalent: easy to predict, expect to know
• Multivalent: product or hint must be given to know
• Non-metal to non-metal
• molecular compound
• Non-metal oxide + water -> acid
• CO2 + H2O -> H2(CO3) (aq)
• P5O5 + H2O -> 2H2(PO3) (aq)
• Choose between PO3 or PO4
• Test which one is balanceable
• Use the polyatomic molecule that best relates to the reactants
• Metal Oxide + Water -> base
• NaO2 + H2O -> 2h2PO4
• Decomposition Reaction
• Binary Compound -> metal + non-metal
• 1) write out result and metal
• 2) Non-metal follows HOFBrINCl method
• 3) Always use original charges
• 4) balance equation
• Metal Nitrate -> metal nitrite + O2
• Metal Carbonate -> metal oxide + CO2
• Combustion Reaction
• Complete Combustion: Cx + Hy + O2 -> CO2 + H2O
• Incomplete Combustion: Cx + Hy + O2 -> CO2 + H2O + CO + C
• Single Displacement Reaction
• Metals displacing metals
• According to the Activity Series, whichever (that’s single) is higher than the one that’s bonding, will do the bonding
• Charges of the metals still apply as they would
• Metals displacing hydrogen
• The hydrogen will be displaced outwards, following HOFBrINCl
• Metals Displacing water
• Only one hydrogen will be displaced outwards
• Results in Metal hydroxide + hydrogen gas
• Charges still apply
• Activity Series of halogens
• Based on the halogen activity series, follows same rules
• Double Displacement Reaction
• Neutralization reaction
• Acid (H) + bases (OH) -> salt + water
• take the positive elements, and switch them
• add water to whatever the salt is, then balance
• charges still apply
• Precipitate Formations
• when it’s a double displacement reaction, just switch the first 2 elements
• AB + CD -> AD + CB
• AgNO3 + NaCl