Chemistry Unit

Atoms, Ions, Molecules and the Periodic Table

Atom

• The basic unit in a chemical element

• Composed of:

  • Protons (nucleus, positively-charged)

  • Neutrons (nucleus, no charge)

  • Electrons (orbits, negatively charged)

Atomic Number and Mass Number

• Atomic number (Z) = to the number of protons in an atom

  • Z = # of protons

• Mass number (A) = total mass of protons and neutrons in an atom

  • A = # of protons + # of neutrons

Protons, Neutrons, Electrons

In a neutral atom, the number of protons will always equal the number of electrons

In Ions

If the charge is positive, that means the atom lost electrons. If it is negative, then the atom gained electrons

Ions

• An atom with a charge due to the gaining or losing of electrons

• Metals tend to lose electrons to become cations

• Non-metals tend to gain electrons to become anions

Terms

Molecule - a group of atoms bonded together

Element - the “type” of atom (one type in a pure substance)

Compound - two or more “types” of atoms bonded together

Periodic Table

• Rows = periods

• columns = groups

• contains information about element symbol, name, atomic number and mass

• Key groups to know:

  • Alkali metals

  • transitional metals

  • Metalloids

  • halogens

  • noble gases

Staircase

• Divides metals and non-metals

• Metalloids are found along the staircase

Trends in the Periodic Table

Down a Group

• The number of valence electrons stays the same

• The number of orbitals increases

• The atomic radius increases

  • Because there is an increasing number of orbits, therefore increasing the size of an atom

Across a Period

• The number of valence electrons increases

• The number of protons increases

• The atomic radius decreases

  • Because there are more protons in the nucleus attaching the electrons

Trends in Reactivity

• The ability of an element to react rests on the atom's ability to lose or gain electrons.

• Metals

  • Metals lose electrons to become positive and reactivity increases when going down the group because the atomic radius increases, meaning weaker attraction from protons.

  • Reactivity decreases across a period because atomic radius decreases, and more valence electrons lose

• Non-metals

  • Non-metals gain electrons to become negatively charged and reactivity decreases because the atomic radius increases, meaning weaker attraction from protons

  • Reactivity increases across a period because the atomic radius decreases and it’s easier to gain electrons.

Trends in the Periodic Table Definitions

Atomic Radius

• The distance measured from the center of the nucleus to the outermost electron level.

• How “big” an atom is

• Correlated with the number of shells (orbitals) and the number of protons in the nucleus (nuclear charge)

• Increases down a group

• Increases across a period from left to right

• Increase in the number of energy levels, therefore electrons are farther away

• Electrons are held more tightly as you move from left to right due to the presence of more protons in the nucleus.

Reactivity Metals

• The degree to which metals have a tendency to react with other substances by losing electrons

• Increases down a group

• Increases across a period from right to left

• Reactivity decreases across a period because atomic radius decreases and more valence electrons are lost

• More energy levels, larger atomic radius, so there is a weaker attraction of electrons

• Weaking attraction of electrons so they are more easily removed

Reactivity Non-Metals

• The degree to which non-metals have a tendency to react with other substances by gaining electrons

• Increase up a group

• Increases across a period from left to right (not including group 18)

• Reactivity increases across a group because atomic radius decreases and it’s easier to gain electrons

• Fewer energy levels, smaller atomic radius, so greater attraction of electrons

• Greater attraction for electrons due to smaller atomic radius

Ionization energy

How much energy is needed to remove an electron from the valence shell?

• Correlated with atomic radius.

• Ionization energy decreases as you move down a group (protons in the nucleus have a weaker pull on valence electrons)

• Ionization energy increases as you move across a period (atomic radius decreases, meaning stronger pull from protons in the nucleus).

• As you move down the group, it becomes easier to remove electrons.

• As you move across a period, it becomes more difficult to remove electrons.

Electronegativity

• How strongly an atom attracts shared electrons in chemical bonds (common in covalent bonds)

  • Electronegativity looks at which one is stronger at attracting shared bonds

• Correlated with atomic radius

• Electronegativity decreases as you move down a group (protons in the nucleus have less attraction on bonding electrons due to larger atomic radius)

• Electronegativity increases as you move across a period (a smaller atomic radius means stronger attraction to bonding electrons)

Fluorine is the

most

electronegative and Francium is the

least

Naming Ionic Compounds

1. The cation will be written first and the name will be unchanged from the element name

2. The anion will be written after the cation and the ending will be changed to “-ide”

   1. E.g., NaCl = sodium chloride

Naming Multivalent Ionic Compounds

1. Write the cation (metal) name first changed and then add a Roman numeral in brackets after to indicate the charge

2. Write the anion name next with the “-ide”

   1. E.g. FeCl2

Naming Ionic Compounds with Polyatomic Ions

• The naming rules for polyatomic ions follow the same rules as you learned previously for ionic compounds

• You simply replace the anion name with the polyatomic ion name

• For example:

  • NaOH → Sodium hydroxide

  • MgSO4 → magnesium sulphate

Naming Molecular Compounds

1. Write the name of the first element in the chemical formula

2. Write the name of the second element and replace the ending with “-ide”

   1. For example: PCl3 → phosphorus chloride

3. Add prefixes depending on the number of atoms for each element

   1. For example: PCl3 →  monophosphorus trichloride

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4. If the first element has a “mono-” prefix, drop the prefix and write the final name

   1. For example: PCl3 → phosphorus trichloride

Chemical equations

Chemical equations is a way to write chemical reactions

Chemical equations can help scientists describe chemical reactions in a standard way

Recall, chemical reaction:

• Changes in the chemical composition of materials

• New substance(s) may be formed

• Difficult to reverse

• This is in contrast to physical changes where the chemical composition of materials stays the same

Parts of a Chemical Equation

1. Reactants

   • Substances before a chemical reaction has occurred; can be 1 or more substances

2. States of matter

   • Written at the end in subscripts with brackets

3. Plus

   • Indicates multiple reactants or products

4. Arrows

   • Indicates a chemical reaction

5. Products

   • Sunstances produced after a chemical reaction; can be 1 or more substances.

Note: This is a chemical change because new substances are produced

States of matter in chemistry

• As you have just seen, chemical equations require states of matter be written

  • G for gas

  • L for liquid

  • S for solid

  • Aq for aqueous

In chemistry, the aqueous state refers to substances dissolved in water. For example, NaCl dissolved in water is an aqueous state.

Arrows in chemical changes

→ Reactions occur in one direction only

⇋ Reactions can occur in both directions (in equilibrium)

Types of Chemical Equations

1. Word equations

   1. A way of describing a chemical reaction using the names of the reactants and products

      1. Water + carbon → glucose + oxygen gas

2. Skeleton equation

   1. A way of describing a chemical reaction using the chemical formulas of the reactants and products (no coefficients are added)

Energy is represented sometimes

Sometimes, energy is needed or produced during chemical reactions, and that can be represented in a chemical equation

Word: Iron + sulphur → iron (III) sulphide + energy

Skeleton: Fe(s) + S(s) → FeS(s) + energy

Energy is produced. Reactants are ion and sulphur solids. Product is iron (II) sulphide.

How to determine states of matter (general rule)

• If it is an element or a diatomic element, please refer to the state on the periodic table at room temperature.

• If it is an ionic compound, please refer to solubility table

• If it is molecular compound, information will be provided on the state of matter

• If it is an acid or base, assume that they are soluable (aq)

How To Read The Solubility Table

Soluable: can dissolve in water

Ions typically have the property of being able to dissolve in water, but some don’t

1. Look for the ions contained within your compound in the first column (i.e. for NaCl, it contains a halide ion)

2. Look for the solubility for your ion in the second column (i.e. halides are soluble)

3. Look for exceptions to the rule in the third column (i.e. halides containing Ag+, Pb2+, Hg2+ and Cu+ are exceptions and are insoluble)

Therefore, based on this information, NaCl is soluble  and exists in an aqueous state in most reactions

When an ionic compound, put (aq), but if it is not soluble, then put (s) for solid

Other things you may see in chemical equations

Coeffieicents

• Show how many molecules are needed in balanced chemical equations; they are not added to skeleton equations

Catalysts

• Usually written in small letters by the arrow; catalysts often help facilitate or speed up chemical reactions, but they themselves are not consumed in these reactions, unlike reactants

Lewis Dot Diagram for Atoms

Steps

1. Write the chemical symbol

2. Find the number of valence electrons

3. Draw the number of valence electrons (recall how you fill electrons in bohr-rutherford diagrams)

Lewis Dot Diagram for Ions

Steps:

1. Write chemical symbol

2. Find the number of valence electrons and draw them

3. Add the number of electrons equal to negative charge or take away electrons to positive charge

4. Draw a square bracket around the ion and write the charge on the outside

Lewis Dot Diagram for Molecular Compounds

1. Draw the Lewis dot diagrams for the atoms

2. Circle pairs of bonding electrons for the formation diagram (share)

   1. No arrows

3. Draw the molecule by replacing the circle electrons with lines (one circle = one line). Each line represents a bond.

   1. Each line represents a bond

   2. One line = single bond, two lines = double bond, etc.

Ionic Compounds

Ionic bonding is the transfer of electrons. Oppositely charged ions attract each other and form an ionic compound bonded with an ionic bond.

Properties of ionic compounds:

• Ionic compounds are solid at room temperature

• Ionic compounds tend to form crystals (crystalline latus)

• Ionic compounds are generally soluble

• Ionic compounds are hard

• Ionic compounds are brittle

• Ionic compounds are crystalline

• Ionic compounds have high melting points

• Ionic compounds are mostly soluble in water

• Ionic compounds dissolve into electrolytes (conduct electricity in water)

  • Dissolves into their ion from (NaCl →Na+Cl-)

  • Pure water is not conductive; only when things have dissolved into the water, releasing ions, will the water be conductive

Forming Ionic Compounds (Zero-Sum)

• Only applicable to ionic compounds; not molecular compounds.

1. Write the symbols of the elements with cations on the left and anions on the right

   1. Mg Cl

2. Add the ionic charge of each ion above the symbol

   1. +2 -1

      Mg Cl

3. Determine how many ions of each type are required to bring the total charge to 0

   1. 1x (+2) 2x(-1) = 0

      Mg Cl

4. Write the chemical formula using the number of ions as subscripts (1 need not be written)

   1. Mg1Cl2 which should be written as MgCl2

   2. Ensure to reduce to the lowest ratio

Forming Ionic compounds (Criss Cross)

1. Write the symbols of the elements with cations on the left and anions on the right

Mg Cl

1. Add the ionic charge of each ion above the symbol

+2 -1

Mg Cl

1. Crisscross the charges and write them as subscripts

   1. Mg1Cl2

2. Divide by the largest common denominator if needed

   1. If it can be reduced, always reduce it

   2. Mg subscripts cannot be reduced

   3. 1:1/1:2 is the lowest ratio

Naming Ionic Compounds

1. The cation will be written first and the name will be unchanged from the element name

2. The anion will be written after the cation and the ending will be changed to “-ide”

   1. E.g., NaCl = sodium chloride

Monovalence

• Elements such as lithium and calcium are known as monovalent elements because they can only have one charge

• Meaning lithium can only lose 1 electron during ionic bonding and calcium can only lose 2.

Multivalence

• Elements such as iron and lead are known as multivalent elements because they can only have multiple charges

• Iron can lose 2 OR 3 electrons during ionic bonding and lead can lose 2 OR 4 electrons

• Typically in the transition medals section

Polyatomic Ions

• Groups of atoms bonded together and treated as one ion

• “Poly-” = many, “atomic” = atoms

• Gypsum contains a polyatomic ion

  • Therefore, $Ca^2$+ is the cation in the compound

  • $SO_4^2$- is the anion in the compound (polyatomic)

  • Together they form an ionic bond

Important note:

• When working with polyatomic ions, you are treating the whole ion as one entity. Therefore, any subscripts apply to the entire ion. Brackets are needed around the polyatomic ion to demonstrate this.

Note on counting atoms:

Similar to math, the subscripts on polyatomic ions can be treated as a multiplication for the entire bracket.

Molecular Compounds

• Forms individual molecules

• Bonded by covalent bonds

• Non-metals

• Very important to biology

• Can be solids, liquids, or gases

• Poor conductors

• Less soluble than ionic compounds

• Molecules can be large

Naming Molecular Compounds

1. Write the name of the first element in the chemical formula

2. Write the name of the second element and replace the ending with “-ide”

   1. For example: PCl3 → phosphorus chloride

3. Add prefixes depending on the number of atoms for each element

   1. For example: PCl3 →  monophosphorus trichloride

   2. If the first element has a “mono-” prefix, drop the prefix and write the final name

      1. For example: PCl3 → phosphorus trichloride

Diatomic Molecules

• Any compound made up of two atoms (could be same or different)

• However, there are only 7 naturally-occurring diatomic molecules formed by two atoms of the same element

• All 7 contain covalent bonds

Balancing Chemical Equations

Balancing chemical reactions

A balanced chemical equation requires the addition of coefficients to the skeleton equation

The coefficients are meant to “balance” the number of atoms for the reactants and products

Example: Hydrogen gas + oxygen gas → water

2 H₂ + O₂ → 2 H₂O

You can ensure that the masses of the products and reactants stay the same when we are writing chemical equation by counting atoms.

Law of Conservation of Mass

Mass cannot be created or destroyed (i.e. atoms cannot be created or destroyed during chemical reactions)

Total mass of reactants equals the total mass of the products

Acid/Base Nomenclature

What are acids and bases?

• Acids are compounds that release hydrogen ions in water when dissolved

• Bases release hydroxide ions (mostly)

• Based on Arrhenius’ definition of acids and bases

• Can be measured using the pH scale

• For acids, the closer the scale is to 0, the stronger the acid

• For bases, the closer the scale is to 14, the stronger the base

Naming acids

• Two main types: binary acids and oxyacids

• Binary acids contain two elements: hydrogen and a halogen

• Oxyacids contain hydrogen ions and another polyatomic ion with oxygen

Naming binary acids

1. Add a “hydro-” prefix

2. Join the anion name with “-ic” ending to the prefix

3. Write the word acid at the end

For example:

HCl hydrochloric acid

HF hydrofluoric acid

When figuring out the formula, the zero-sum and criss-cross methods still work.

Naming oxyacids

The name of the oxyacid depends on the polyatomic ions bonded to the hydrogen ion(s)

Naming bases:

Most bases contain hydroxide ions, and you name them similarly to how you named polyatomic ions

For example

• NaOH = Sodium hydroxide

• LiOH = Lithium hydroxide

However, not all bases contain hydroxide ions, such as ammonia (NH3)

Additional note:

• Acid and bases are not separate categories but are very interconnected concepts

• During a chemical reaction, an acid usually reacts to form a conjugate base (i.e. an acid-base pair)

• Water can be considered as an acid and a base as it self-ionizes (forms ions from itself)

Types of Reactions: Synthesis and Decomposition

Synthesis Reaction

• Two or more reactants form one product

• General equation: A + B → AB

  • A can be a compound, B can be a compound; it can be compound + compound forming another compound

  • There can be more than two reactants; as long as it forms one product, it is considered a Synthesis reaction

• Examples:

  • Hydrogen gas + oxygen gas → water

  • Carbon + Oxygen → carbon dioxide

General Rules and Examples (Do not have to memorize)

1. Element + element → binary compound

   1. 2Na + Cl2 → 2NaCl

2. Metal oxide + CO2 → Metal carbonate

   1. MgO + CO2 → MgCO3

3. Metal + Oxygen gas → metal oxide

   1. 4 Na + O₂ → 2 Na₂O

4. Metal oxide + H₂O → base (hydroxide)

   1. MgO + H₂O → Mg(OH)₂

5. Non-metal oxide + H₂O → oxyacid

   1. CO₂ + H₂O→ H₂CO₃

Special Synthesis Reactions (memorize)

If asked to predict a synthesis reaction, it will be one of the following:

• 2 H2 (g) + O2 (g) → 2 H2O (l)

• 2 H2O (l) + O2 (g) → 2 H2O2 (l)

• C (s) + O2 (g) → CO2 (g)

• 2 CO (g) + O2 (g) → 2 CO2 (g)

• N2 (g) + 3 H2 (g) → 2 NH3 (g)

• H2O (l) + CO2 (g) → H2CO3 (aq)

Decomposition Reaction

• One reactant breaks down to form two or more products

• General equation: AB → A + B

  • The opposite of Synthesis’s reaction

  • A and B don’t have to be elements; they may be compounds

Decomposition Rules

1. Binary compound → element + element

2. Metal carbonate → metal oxide + CO₂

3. Metal oxide → metal + O₂

4. Base (hydroxide) → metal oxide + H₂O

5. Oxyacid → non-metal oxide + H₂O

Special Decomposition Reactions

If asked to predict a decomposition reaction, it will be one of the following:

• 2 H2O (l) → 2 H2 (g) + O2 (g)

• 2 H2O2 (l) → 2 H2O (l) + O2 (g)

• CO2 (g) → C (s) + O2 (g)

• 2 CO2 (g) → 2 CO (g) + O2 (g)

• 2 NH3 (g) → N2 (g) + 3 H2 (g)

• H2CO3 (aq) → H2O (l) + CO2 (g)

Types of Reactions: Single Displacement

Single Displacement Reactions

• Element (or diatomic molecule) reacts with a compound to form a different compound and element (or diatomic molecule)

• General equation: A + BC → AC + B

Metal Activity Series

• Lists metals based on reactivity

• In a single displacement reaction, more reactive metals will displace less reactive metals from the compound

• A less reactive metal will NEVER replace a more reactive metal in a single displacement reaction

• No reaction will occur as a result

• Represented by writing NR on the products side

  • Ex. Ca + KCl → NR

Types of Reactions: Double Displacement

Double Displacement Reactions

Two compounds react to form two different compounds

General equation: AB + CD → AD + CB

Note: Order matters. A and C are typically cations and will still be cations in the new reaction, while B and D are anions.

Examples:

Ions will NOT change their charges during double displacement reactions. If copper has a charge of 1+ in the reactants side, it will also have a charge of 1+ in the products side.

Precipitate

A common sign of a double displacement reaction is the formation of a solid product known as a precipitate. Reactants of double displacement reactions are usually aqueous solutions.

Predicting Double Displacement Reactions

With double displacement reactions, we will just assume that there is always a reaction.

To predict the double displacement reactions,

Identify the cations and anions in each reactant.

Switch the pairings between the cations and anions

Write the products with their states

Balance out the equation

Neutralization

A special type of double displacement: A reaction between an acid and a base

General reaction: acid + base → water + salt

Neutralization reactions will always form water and a salt (salt in chemistry is NOT referring to table salt but rather an ionic compound)

Types of Reactions: Combustion

• Occurs when a fuel reacts with oxygen to produce oxide and energy

• Mainly focusing on hydrocarbons

  • As the name suggests, hydrocarbons contain hydrogen and carbon

  • Methane (CH₄) is an example of a hydrocarbon

Two types of Combustion

Complete combustion: occurs in the presence of adequate oxygen

General reaction: $C_HH_y+O_2(g)→ CO_2+H_2O(l)+ energy$

Incomplete combustion: Occurs in the presence of inadequate oxygen

General equation: $C_HH_y+O_2(g)→ CO_2+H_2O(l)+CO(g)+C(s)+ energy$

Complete Combustion

• Occurs in the presence of adequate oxygen (efficient energy production)

• Bunsen burners are common lab equipment used and involve the combustion of methane

• Complete combustion of hydrocarbons will usually yield a blue flame that burns very hot

Incomplete Combustion

• Occurs in the presence of inadequate oxygen (inefficient energy production)

• Incomplete combustion of hydrocarbons will usually yield yellow flame, soot, and dangerous CO gas

Acids, Bases, and the pH Scale

Arrhenius Acids and Bases

Acid - Forms H in water

HCl → H + Cl

Base - forms OH in water

NaOH → Na + OH

Properties of Acids

• Release H in water

• React with metals and carbonates

• Conduct electricity (electrolytes)

• Taste sour (never do this)

• Corrosive

• Neutralize bases

Properties of Bases

• Release OH in water

• Conduct electricity (electrolytes)

• Taste bitter (never do this)

• Feels slippery

• Corrosive

• Neutralize acids

pH Scale

• pH stands for the “Potential (or power) of Hydrogen”

• Measured on a scale of 0-14

• If the pH of a substance <7, it is considered acidic

• If the pH of a substance >7, it is considered basic (or alkaline)

• If the pH of a substance = 7, it is considered neutral

Potential of Hydrogen

• The pH scale actually indicates the concentration of hydrogen ions in a solution

• The stronger the acid, the more concentration of hydrogen

  • Note: concentration of H will be denoted as [H*]

  • Square brackets refer to the concentration of something in a solution

• Calculated as follows:

  • pH = -log[H+]

    • Will not be asked to calculate

It is a logarithmic scale, meaning every time you move it, it is 100

Think of exponents in math.

Strong and Weak Acids and Bases

• Strong acids and bases - completely dissociate in water to form ions (i.e., all molecules will break apart into ions)

• Weak acids and bases - partially dissociate in water to form ions (i.e., some molecules will break apart into ions)

• A strong acid will be more acidic than a weak acid (more H* released)

• A strong base will be more basic than a weak base (more OH released)

Measuring pH

It is impossible to tell the pH of something by looking at it. There are three methods used to identify the pH of a solution:

1. Litmus test

2. pH indicator

3. pH probe

Litmus paper (red/blue)

Good for testing if something is acidic or basic, but it cannot tell you what the exact pH it is.

Litmus paper (universal)

Displays a range of colours depending on the pH of the substance

pH Indicators

• Chemicals commonly used in titrations to track changes in pH

• Phenolphthalein is a commonly used indicator for acid-base neutralization reactions because it changes colour at neutral

• Also, know menthyl orange

• Some indicators may change colours more than once

pH Probes

• Very accurate reading of pH

• Very expensive

• Can also track changes in pH

• Can link to computers to store data