1 – Properties of Matter

Science 10 – Unit 2 – Chemical Reactions

The Particle Theory

The particle theory is used to explain many observable characteristics and properties of different substances.

Before learning the particle theory, it is important to know the definition of matter:

Matter is anything that has mass and takes up space (volume).

It can be solid, liquid, gas or plasma.

The Particle Theory

The particle theory has four main tenets:

1. All matter is made up of particles.

2. Particles are always in motion, since they contain energy.

3. There is empty space (nothing!) in between particles.

4. The particles of one substance have different characteristics than particles of other substances.

Particle Theory

Note that “particles” is a non-specific word. It can be used to mean:

Atoms, which are particles of an element;

OR

Molecules, which are particles of a compound;

OR

Ions, which are particles that have a positive or negative charge.

Classifying Substances

• Elements are made up of one type of atom only, and can be found on the periodic table

• Compounds are made up of more than one type of atom, combined into all of the same kind of molecules, but cannot be separated

• Mixtures contain more than one type of compound or element and can be separated into the individual components

Definition of Physical and Chemical Properties

Physical properties are characteristics of a substance that can be observed without changing its chemical composition. Physical properties also include information on how a substance changes state.

Chemical properties describe how a substance reacts when it interacts with another substance.

Qualitative and Quantitative Properties

Properties of matter can sometimes be subjective, meaning the actual observation might vary depending on the observer or the sample. Other times, properties may be exact and consistent for multiple samples and observers.

Qualitative properties are characteristics that can only be observed, not measured.

Quantitative properties are measurable.

Chemical and Physical Properties

Physical Properties:

Colour

Lustre (shininess)

Melting point

Boiling point

Density

Solubility

Ductility

Conductivity

Hardness

Texture

Malleability

Viscosity

Odour

Brittleness

Chemical and Physical Properties

Other physical properties can include crystal shape, state, size, shape, volume, temperature and mass.

When describing physical properties of a substance, it is best to use as much detail as possible.

Chemical and Physical Properties

Chemical Properties:

Reactivity

Flammability

Toxicity

Heat of reaction (energy released/absorbed)

Chemical stability

Physical and Chemical Changes

• A physical change occurs when a substance is altered in terms of its shape, size, state, etc. However, in a physical change, the composition of the substance does not change.

• A chemical change occurs when a substance is transformed into a new, different substance

Signs of a Chemical Change

• A new colour appears

• Heat or light is given off

• Gas is produced (bubbles)

• A solid is produced (called a precipitate)

• A temperature change occurs

Remember: the main sign of a chemical change is that a new substance is formed!

2 – The Periodic Table

Science 10 – Unit 2 – Chemical Reactions

Development of the Periodic Table

In 1867, Dmitri Mendeleev began to collect information on all of the known elements at the time (63 of them) and started to sort the elements according to common properties: reactivity, colour, state, mass and density.

As he organized, he began to identify a pattern. He hypothesized that gaps in the table represented elements that had not yet been identified yet, but whose properties could be predicted.

Development of the Periodic Table

Mendeleev’s work led to the periodic law:

When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical

properties.

This means that the periodic table isn’t arranged randomly – it is very logical!

Classes of Elements

Metals are located on the left side of the periodic table.

General Properties of Metals

High lustre (shiny)

Malleable, ductile

Good conductors of heat and electricity

Solid at room temperature (except mercury)

Classes of Elements

Non-metals are located on the right side of the periodic table.

General Properties of Non-Metals

Low lustre (dull)

Brittle

Good insulators of heat and electricity

Some solids, many gases, one liquid

Classes of Elements

Semi-metals, or metalloids, are located to the right of the middle of the periodic table, along the edge of the “staircase”.

Semi-metals have properties of both metals and non-metals. One of the most common uses of semi-metals is in electronics, because they can conduct electricity in a controlled way.

Structure of the Periodic Table

A period is a horizontal row on the periodic table. There are seven periods of elements.

A group is a vertical column on the periodic table. It is also called a chemical family, because the elements in it have similar characteristics.

Families of Elements

There are four specially-named families.

Alkali metals is the first group. They are soft, shiny metals that are very reactive.

Alkaline earth metals is the second group. They are shiny, silvery metals that are somewhat reactive, but not as reactive as alkali metals.

Families of Elements

Halogens are the second last group. They are highly reactive elements that are harmful to living things.

Noble gases are the final group. These gases are all very stable and do not react.

Families of Elements

Transition metals make up the centre block of the periodic table, from the 3rd to the 12th column.

Lanthanides and actinides are in the separate two rows below the main part of the periodic table.

Elements on the Periodic Table

Each block on the periodic table gives information about a specific element. On a standard table, there will be the element symbol, the element name, the atomic number and the atomic mass.

Depending on the complexity of the periodic table, there may be additional information.

Atomic Number

The elements are arranged on the periodic table in order of increasing atomic number.

The atomic number of an element is the number of protons in the nucleus of the atom.

Remember that for an atom, the number of electrons is equal to the number of protons!

Atomic Mass

The atomic mass is the mass of one atom of an element. It is equal to the number of protons and neutrons added together.

The units for atomic mass are amu, which stands for “atomic mass units”.

Atomic Mass

On the periodic table, the atomic mass is actually an average of all of the masses of the naturally-occurring isotopes of a particular elements.

Isotopes are atoms with the same number of protons, but different numbers of neutrons.

Element Symbols

Element symbols are vital for writing the names of compounds and in chemical reactions.

3 – Atomic Structure

Science 10 – Unit 2 – Chemical Reactions

Current Atomic Theory

The particle theory, which we looked at earlier in the unit, is actually the currently accepted atomic theory.

We know that all matter is composed of atoms. Also:

• Atoms of a specific element (like helium) are identical in size and mass.

• Atoms of different elements can combine together to make new substances.

Subatomic Particles

Atoms are not technically the smallest piece of matter. Subatomic particles are the components that make up an atom. Each has a specific purpose in the atom and in chemistry (and other sciences).

There are three main subatomic particles: electrons, protons and neutrons.

Subatomic Particles

Particle Mass Charge Location

Electron Very light Negative Electron cloud

Proton Heavy (~2000 times heavier than

an electron)

Positive Nucleus

Neutron Heavy (slightly heavier than a

proton)

Neutral (no charge)

Nucleus

The Atomic Model

In the current model of the atom, it is known that the protons and neutrons are in a tight mass in the centre of the atom, called the nucleus.

The electrons orbit in a (relatively) huge cloud around the nucleus, in different energy levels based on the number of electrons.

The Atomic Model

Atomic Number

• Recall that the atomic number of an element tells how many protons it has in its nucleus.

• For an atom (which has no charge), the number of electrons and protons are equal.

• Ions are atoms that have gained or lost electrons to give them a charge.

Electrons

• The placement and number of electrons in an atom will be very important as we start to look at chemical bonds and reactions

• There are two ways to “draw” atoms and ions that give important details about the elements

Bohr-Rutherford Diagrams

• The theory put forth by Niels Bohr and Ernest Rutherford claimed that electrons orbited in “shells” around the nucleus of the atom.

• Each shell could hold a specific number of electrons:

Shell 1 2 3 4

Number of Electrons 2 8 8 18

Drawing Bohr-Rutherford Diagrams

As an example, use Mg.

1. Write the chemical symbol in the centre.

2. Draw a circle around the “nucleus”.

3. Draw the first shell, and put two electrons in it on opposite sides (10 left).

4. Add the second shell and draw 8 electrons in pairs (2 left).

5. Add the third shell and start with one electron at the top and one on the right.

Practice

• Draw the Bohr-Rutherford diagrams for each of the first 20 elements on the worksheet.

• Then, answer the questions at the bottom. We will discuss this in 15 minutes.

Lewis Diagrams

• Lewis diagrams only show the chemical symbol and the valence electrons.

• Valence electrons are those in the outermost shell (and available for chemical bonds).

• The electrons are shown as dots starting on the right side and moving clockwise around the atom until it is filled with eight electrons.

• The octet rule states that most elements or molecules are “happy” with eight electrons in their valence shell (one notable exception is sulfur)

Lewis Diagram Examples

1. H

2. Cl

3. Ne

Ions

• Ions are formed when an atom gains or loses electrons

• Non-metals (left side of the PT) gain electrons and become negatively charged

• Metals (right side of the PT) lose electrons and become positively charged

Ion Notation

• When an atom becomes an ion, it is written with the same symbol and the charge (how many electrons it gained or lost) as a superscript

• Example:

Na+ (lost one electron)

N3- (gained three electrons)

Ion Rules

Rules for ions:

–Hydrogen can be ±1 (usually +1)

–Group 1 is 1+, group 2 is 2+, group 3 is 3+

–Transition metals are multivalent (can have different charges)

–Group 15 is 3-, group 16 is 2-, group 17 is 1-

–Group 18 (noble gases) do not become ions

4 – Ionic Bonds

Science 10 – Unit 2 – Chemical Reactions

Ions

• Remember that an ion is an atom that has lost or gained electrons so it has a full valence shell.

• An anion is a negative ion (made from non-metals).

• A cation is a positive ion (made from metals).

Ionic Bonds

• An ionic bond is a bond between two or more ions, with at least one anion and one cation

• The way to identify a bond as an ionic bond is that it contains a metal and at least one non-metal.

Formulas of Ionic Compounds

• Overall charge must be zero (add all of the + and -)

• Metal is always listed first

• Number of ions of each element is indicated by a subscript number

• No charges shown

• Always be VERY careful about capital and lowercase letters!

Nomenclature

• Nomenclature just means “naming”

• Ionic compounds are easy!

• For a “regular” metal and non-metal:

– Positive ion first (metal) – write the full element name

– Negative ion second (non-metal) – first syllable of the element name plus –ide (e.g. oxide, nitride, iodide)

Multivalent Ions

• Some metals have more than one possible charge – these are called multivalent.

• Most are transition metals, in the center block of the periodic table

• The charge on these is indicated by Roman numerals

– e.g. Manganese (II), Manganese (III)

• Some have special names to indicate which charge they have:

– e.g. Iron (II) is ferrous, iron (III) is ferric

Polyatomic Ions

• Polyatomic ions are groups of atoms that have an overall charge.

• They can be considered a “unit” – all together!

• Examples:

– OH- = hydroxide

– CO32- = carbonate

• When you are writing them in a compound, they need to have brackets if there is more than one:

– e.g. KOH, Ca(OH)2

Review So Far…

• Elements vs compounds vs mixtures

• Structure of the periodic table

– Metals, non-metals and metalloids

– Special families

– Using the PT to determine structure of an atom

• Atomic diagrams and valence electrons

– Bohr-Rutherford diagrams

– Lewis diagrams

Review So Far…

• Ions

– What are they? How are they made?

– “Special” ions (multivalent and polyatomic)

• Ionic bonds

– Characteristics

– Naming

– Writing formulas

5 – Covalent Bonds

Science 10 – Unit 2 – Chemical Reactions

Ionic versus Covalent Bonds

Ionic Covalent (Molecular)

• Metal + non-metal(s) • Electrons are

transferred from metal to non-metal

• Made of ions in a lattice structure

• Only non-metals • Electrons are shared

between atoms • Made of molecules

Covalent Compounds

• As previously mentioned, covalent compounds are those that only contain non-metals.

• You only need to know how to write names and formulas for binary compounds, which are those that only have two elements in them. It gets much more complicated when there are more!

Naming Covalent Compounds

• Naming covalent compounds is almost the same as naming ionic compounds, except that covalent compounds have prefixes

• The first element only has a prefix if there is more than one of them

• The second element always has a prefix and ends in “-ide”

Prefixes

mono- 1

di- 2

tri- 3

tetra- 4

penta- 5

hexa- 6

hepta- 7

octa- 8

nona- 9

deca- 10

Common Names and Exceptions

• HF, HCl, HBr, HI and H2S never have prefixes

• You can name common compounds with the common name

Water H2O

Ammonia NH3

Methane CH4

Propane C3H8

Ethane C2H6

Examples

• Cl3F

• N2O5

• CO

• Cl2O7

• CCl4

• HCl

• H2O

Writing Covalent Formulas

• Using the name, write the element symbols in the same order.

• Use the prefixes to determine the number of each atom. Remember that if there is no prefix, there is only one of the atom. The only exceptions to this are hydrogen sulfide (H2S) and the diatomic gases.

Diatomic Gases

• There are seven diatomic elements (ones that come only in groups of two).

• Diatomic: di = two, atomic = atoms

• They are:

H2, N2, F2, O2, I2, Cl2 and Br2

• Sometimes, sulfur and phosphorus (as elements) are written as S8 and P4.

Examples

• carbon dioxide

• ammonia

• silicon tetrachloride

• diphosphorus pentasulfide

• tetrachlorine hexafluoride

• nitrogen gas

• hydrogen sulfide

6 – Writing and Balancing Chemical Equations

Science 10 – Unit 2 – Chemical Reactions

Chemical Equations

• A chemical equation is a way of writing out how a chemical reaction will occur, using the formulas for each of the compounds that participate.

• A proper chemical equation will have all of the compounds and elements, their states and will be properly balanced (more on that later)

General Form aA + bB → cC + dD

• Each of A, B, C and D is a chemical compound (e.g. H2O)

• Each of a, b, c and d is a number called a coefficient that says how many molecules of that compound are needed for the reaction to happen

• + means “and” and → means “reacts to form”

• A and B are reactants, C and D are products

Remember!

• Remember that a chemical change (or a reaction) occurs when a new substance or substances are formed.

• Now we know these are called products, and we can actually figure out the compounds that will be made when two things are mixed together.

Example Reaction

Example Reaction

2 HCl (aq) + Mg (s) → MgCl2 (aq) + H2 (g)

• This reaction also includes states.

• The states for reactions are:

– s for solid

– l for liquid

– g for gas

– aq means “aqueous” which refers to a substances that is dissolved in water

Writing Chemical Reactions

• Reactions can be written using a word equation:

sodium + water → sodium hydroxide + hydrogen

• Or they can be written using the chemical formulas:

Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g)

Example

Write as a word equation and using chemical formulas:

Lead nitrate solution is mixed with potassium iodide solution, which creates a lead iodide solid and a potassium nitrate

solution.

Example

For the previous reaction:

1. What are the reactants?

2. What are the products?

3. What other states are there that are not mentioned in the reaction?

4. Read the reaction out loud!

Practice

• pg. 219 “Try This Activity”

• pg. 219 #1-6

Law of Conservation of Mass

• The Law of Conservation of Mass states that, in a chemical reaction, the mass of the reactants will be equal to the mass of the products.

• This is because the number and type of atoms are the same for the reactants and products, just rearranged.

Balancing Equations

• Balancing an equation means identifying how many molecules of each compound are used up as reactants and created as products.

• For example:

N2 (g) + 3 H2 (g) → 2 NH3 (g)

One molecule of nitrogen combines with three molecules of hydrogen to make two ammonia molecules.

Balancing Equations

• Remember that the numbers added into a chemical equation to balance it are called coefficients.

• Once the coefficients are added, the number of atoms of each element should be the same on each side.

How to Balance

1. Balance the metals.

2. Balance the non-metals that are not hydrogen or oxygen.

3. Balance hydrogen and oxygen.

4. After each step, recheck the previous steps and be sure to double check all numbers at the end!

Example 1

___ KCl → ___ K + ___ Cl2

Example 2

___ H3PO4 (aq) + ___ NaOH(aq) → ___ Na3PO4(aq) + ___ H2O(l)

Example 3

___ Ca(AlO2)2 + ___ HCl → ___ AlCl3 + ___ CaCl2 + ___H2O

Example 4

___ C2H6 + ___ O2 → ___ CO2 + ___H2O

7 – Types of Chemical Reactions

Science 10 – Unit 2 – Chemical Reactions

Types of Chemical Reactions

• There are five main categories of chemical reactions.

• Each category has more specific sub-categories, but in general, all reactions will fit into one main category.

• Once you can identify types of reactions, it lets you predict the products that will form if two reactants are mixed together.

Synthesis Reactions

• Two simple compounds combine to make one more complex compound:

A + B → AB

• An example of this type of reaction occurs when solid potassium is put in a container filled with chlorine gas, and potassium chloride is formed.

• You can identify this reaction by its one product.

Decomposition Reactions

• One compound breaks apart to make two (or more) simpler compounds:

AB → A + B

• An example of this type of reaction occurs when calcium carbonate is heated, causing it to decompose into carbon dioxide and calcium oxide.

• You can identify this reaction by its one reactant.

Single Displacement Reactions

• An element “switches places” with an atom that is part of a compound:

A + BC → B + AC

• When copper metal is put into silver nitrate solution, the copper replaces silver to make solid silver and copper nitrate solution.

• You can identify this reaction because it has a single element on both sides of the reaction.

Double Displacement Reactions

• An atom that is part of a compound “switches places” with an atom that is part of another compound:

AB + CD → AD + CB

• AgNO3 + NaCl → NaNO3 + AgCl

• You can identify this reaction because it has two compounds on both sides of the reaction, all of which are ionic.

Combustion Reactions

• A combustion reaction occurs when a hydrocarbon (an organic compound containing mostly hydrogen and carbon) reacts with oxygen.

CxHy + O2 → CO2 + H2O

• You can identify this reaction because it always contains oxygen gas as a reactant and carbon dioxide and water as the only products.