SCH3U0

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Mltiple intellegence - Concentration

Ved Hailey S. Ingram

Percents

Ved Arianne Boynton

Measurement (length, area, and volume)

Ved Anna Rejda

Chemical Energy

Ved John Doe

Floating topic

Combined Gas Law: (P1V1)/T1 = (P2V2)/T2

SCH3U0

Solutions & Solubility

Stoich and Solutions

Expanding to include concentrations and volume into regular stoichiometry problems

Formulas:

n = m/M

c1v1 = c2v2

C = n/v

Chemical Equations

3) Net Ionic Equation

'Spectator Ions' removed Spectator Ions are ions that don't appear on both sides of the equation

2) Ionic Equation

Each Ion separated

1) Balanced chemical equation

Creating Solutions

Two liquids mix, one solid (precipitate)

Dilutions

Dilution Formula: C1V1 = C2V2

Diluting a stock solution of HCL(aq)

- Add 250mL or more of distilled water up to the fill line

- Measure and add 250 mL of HCL (aq) to volumetric flask

- Add approx. 500 mL of distilled water to volumetric flask

- Obtain 1000mL volumetric flask

From a solid

Creating 500 mL of a mol/L of NaOH

- Obtain 500mL volumetric flask

- Add 250 mL or more to reach the fill line

- Cap and invert to mix

- add approx. 250 mL of distilled water to flask

- Place measured amount of solid in flask

Concentration

How much of solute in a solution in terms of moles per litre(mol/L) or "M".

Molar Concentration Formula

Concentration = # of moles/volume or (c = n/v)

Percentage Concentrations

Other Concentrations

ppm (parts per million) = - 1 g/m3 - 1 g/1000L - 1 mg/L - 1 mg/kg

Used for various purposes in industries that use different standards

3 types

Mass/Volume % (m/v)

Volume/Volume % (v/v)

Mass/Mass % (m/m)

Components of a solution

Solvent

Medium in which solute is being dissolved (water, alcohol, etc.)

Solute

Substance being dissolved into a solvent (sugar, salt, etc.)

Characteristics of solutions

Particles cannot be clearly seen

Cannot be filtered

Unsaturated Solution - Solute less than fully saturated

Saturated Solution - No more solute can be dissolved at the given temperature

Polar molecules dissolve in polar non-polar dissolve in non-polar

May be slightly coloured

Transparent (Clear if aqueous)

Homogenous Mixture

- Can be separated into components - Looks the same throughout

Gases & Atmospheric Chemistry

Dalton's Law of Partial Pressure

The pressure that each gas in a mixture exerts is it's partial pressure

Partial Pressure is used to determine the concentration or fraction in a mixture

P(total) = P1 + P2 + P3...

Gas Stoich

Expanding to include pressure and temperature into regular stoichiometry problems.

Formulas needed:

The "Ideal" Gas Law

SATP is Standard Ambient Temperature and Pressure --------------- SATP = 298K and 100 kPa

STP is Standard Temperature and Pressure STP = 273K and 101.3 kPa

Pressure multiplied by volume is equal to the number of moles times the universal gas constant and temperature

PV = nRT

R = 8.31 kPa*L/mol*K

Gay-Lussac Law

The pressure of a gas varies directly with the temperature of the gas. (At a constant volume)

P1/T1 = P2/T2

Charles' Law

Temperature

Absolute Zero: Theoretical Temperature where all matter would freeze. Absolute Zero is "-273 C" or "0 K"

Temperature recorded in 3 ways: Celsius (C) Fahrenheit (F) Kelvin (K) K = C + 273

The more kinetic energy the particles have, the higher the temperature

The volume of a fixed amount of gas varies directly with the temperature of the gas (constant pressure)

V1/T1 = V2/T2

Calculations always done in Kelvin

Boyle's Law

The volume of a gas varies inversely with with pressure (at a constant temp)

Shampoo bottle exploded on plane - Altitude increases, pressure decreases, making the volume increase

P1V1 = P2V2

Kinetic Molecular Theory

Pressure = force/area

Can be measured in: Kilopascals (kPa) Pascals (Pa) Atmospheres (atm) Millimeters of Mercury (mmHg) Torrs

Gas particles are in constant motion and have perfectly elastic collisions

(in closed systems) Multiple collisions increase pressure

Background/Intro

Gasses are the only compressible state

States of substances depend on the forces between the particles

If the forces are strong - solid If the forces are weak - Liquid or gas

Chemical Trends & Bonding

Bonding

Polar Covalent Bond

For a molecule to be polar, it must be: - Containing a polar bond - Molecule does not have symmetry

An uneven distribution of charges due to unequal sharing of electrons

Intermolecular

Ionic Bonds

Intramolecular

Covalent Bonds

1. Hydrogen bonding 2. Dipole-dipole 3. London forces

Bonding Capacity

Groups (4-8) = 8 - # of valence electrons

Groups (1-3) = # of group

related to # of valence electrons

Lewis-Dot Bonding/Structures

Periodic Trends

PERIODIC LAW: When elements are arranged by increasing atomic number, they fall into categories with similar chemical/physical properties

Atomic Radius

Increases 🡳 and 🡰

Factors affecting size

Strength of attraction of valence electrons

Shielding effect - multiple energy levels

# of shells/energy levels

Half of distance from nucleus to valence

Stoichiometry

Stoich Problems

Molar Ratio: Moles of unknown/moles of known

Percent Yield

Tells us how 'efficient' a reaction was

Actual Yield: Actual amount of product formed

Theoretical Yield: Calculated or 'expected' amount of product formed

% Yield = Actual/Theoretical

Limiting/Excess Reagent

To determine which reactant will be consumed first

Divide moles of each reactant by coefficient of reactant to find # of "cycles"

Reactant with more cycles is excess reactant with less is limiting

Given both reactants' masses

Find mass of either reactant(s) or products

Use formulas: (n = m/M) and (n = N/Na)

Empirical & Molecular Formula

Molecular

Actual ratio for formula of a compound

Empirical or 'Simplest'

Formula for a compound in the smallest whole number ratio

Percent Composition (%)

% composition = (Mass of individual element/total mass) x 100

Moles

Formulas

# of moles = # of entities/Avogadro's constant or n = N/Na

# of moles = mass(g)/Molar mass(g/mol) or n=m/M

1 Mole (or "mol) is ~6.02 x 10^23

**Avogadro's Constant**

*The unit used to measure amounts of any substance.

Chemical Reactions

Nomenclature

Polyatomics and diatomic molecules

Balancing Chemical Equations

Set of rules for naming compounds

Types of Reactions

Combustion

- Burning of hydrocarbons - Complete combustion produces CO2 and H2O

Double Displacement

Neutralization Reactions

Base + Acid 🡲 Salt + Water

Precipitation Reaction

Always has a solid and an aqueous product

Single Displacement

- Element reacts with a compound and displaces the 2nd element to form a new compound and an element

Decomposition

- Reactant broken down into 2 or more products

Synthesis

- 2 different molecules/atoms combine to produce a compound