Science 1206Unit 3 – Chemical Reactions

Text: Chapters 5-8

Chemistry: the study of matter; its properties and its changes.

Matter: anything that has mass and takes up space (energy is not matter).Three states of matter: solid, liquid and gas.

Lab Safety:

Workplace Hazardous Materials Information System (WHMIS):

Three parts: Labeling Worker education and training MSDS (Material safety data sheets)

WHMIS Safety Symbols – labeling (see p.658):

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MSDS – Material Safety Data Sheets:

Provides workers and emergency personnel with the proper procedures for handling and working with a particular substance.

Classification of Matter

Classification of matter as pure substances or mixtures:

Matter can be classified into 2 groups:

1. Pure substances - have constant composition; all the particles that make up the

substance are the same.- Two divisions:

a) Elements: The simplest form of matter that can exist under normal

conditions. Composed of only one kind of atom. Eg. Carbon, magnesium Cannot be broken into simpler substances by chemical means Combine to form other substances

b) Compounds: Substances composed of 2 or more different kinds of atoms. Eg

H2O Can be broken down into simpler substances by chemical means

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2. Mixtures- Have variable composition; composed of 2 or more pure substances.- Two groups:

a) Homogeneous mixture (solutions) Have only1 visible component/phase eg. Tap water. Air, sugar solution (sugar + water)

b) Heterogeneous mixture (mechanical mixture): Have 2 or more visible components Eg, sand in water, vegetable soup, tossed salad.

Properties of Matter – 2 Types

1. Physical Property

Characteristics of matter, used to identify substances Eg. State at room temperature, boiling point (found on periodic table),

color, mass, electrical conductivity

2. Chemical Property Characteristic of matter that can be observed when matter undergoes a

change in composition; Describes “how it reacts” Eg. Propane reacts with oxygen to produce carbon dioxide and water

Changes of Matter – 2 Types

1. Physical Change : A change in the size or form of a substance that does not change its

composition. Eg. Cutting paper, melting ice.

2. Chemical Change: A chemical reaction; a change in which at least 1 or more “new”

substances (products) are formed. The products have different properties from the starting substances

(reactants) Eg. Rust, burning, cooking

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Example: Rusting

Chemical equation for rusting:

Fe(s) + O2(g) ® Fe2O3(s)

The rust produced has completely different properties from iron and oxygen.

Evidence of Chemical Change

Change in color Energy produced (heat and/or light) Gas produced (bubbles) New solid produced (called a precipitate) Difficult to reverse

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Elements and the Periodic Table

All elements are classifies as metals or nonmetals, depending on their properties.

Property Metals Nonmetals

Lustre shiny DullMalleability Malleable (bendable) Brittle

Conductivity of heat and electricity

Good conductors(electrolyte)

Poor or nonconductors(non-electrolyte)

State at room temperature All solids except mercury (liquid)

Most are gases, some are solids and bromine is a

liquidReactivity with acid Mostly yes no

Location (periodic table) Left of staircase lone Right of staircase line

Metalloids (semimetals): Elements that have some properties of metals and some of nonmetals. Include all elements on either side of the staircase line except AL and At. Also includes one form of C, called graphite, which is dull and brittle (nonmetal)

but is a good conductor of electricity (metal).

Chemical Families (groups)

Groups of elements in the same vertical column that have similar physical and chemical properties.

Group 1 (IA) – alkali metals Group 2 (IIA) – alkaline earth metals Group 17 (VIIA) – halogens Group 18 (VIIIA) – noble gases Rows on the bottom:

Lanthanide series Actinide series

Properties of chemical families:Alkali metals:

- group 1, IA- Eg. Sodium, lithium- Show metallic properties- Highly reactive, especially with water; reactivity increases down the group- Cs and Fr are the most reactive metals.

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Alkaline earth metals:- Group 2, IIA- Eg. Calcium, magnesium- show metallic properties- less reactive than alkali metals; reactivity increases down the group

Halogens:- Group 17, VIIA- Eg. Chlorine, fluorine- show nonmetallic properties- reactivity decreases going down the group; F is the most reactive nonmetal- react with metals to produce salts (ionic compounds)- react with hydrogen to form compounds that dissolve in water to form acids.

Noble Gases:- Group 18, VIIIA- Helium, neon, argon- show nonmetallic properties- extremely low chemical reactivity

Hydrogen:- the lightest and most abundant element- doesn’t really belong to any group (can be group 1 or 17)- it sometimes behaves like an alkali metal (group 1) and at other times as an acid.

The ATOM- the basic building block of all matter- the smallest particle of an element that retains the properties of that element- electrically neutral: the number of positive charges (protons) equals the number of

negative charges (electrons)- composed of 3 subatomic particles: protons (p+), electrons (e-) and neutrons (no)

Particle Symbol Relative charge

Mass (g) Location

Proton p+ 1+ 1.67 x 10-24 NucleusNeutron e- 0 1.67 x 10-24 Nucleuselectron no 1- 9.11 x 10-28 orbital

Models of the atom – see handout

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a. Atomic nucleus is located at the core of an atom and it contains protons and neutrons.

[i] Proton: A subatomic particle that carries a positive one (+1) electrical charge and has a mass value of 1 a.m.u. (atomic mass units)

The number of protons in the atomic nucleus determines an element's Atomic Number.

[ii] Neutrons:A subatomic particle that has no electrical charge (0) and has a mass value of 1 a.m.u.

Atomic Mass is the sum of the number of protons and neutrons found in an atom of any element.

b. Energy levels are regions of subatomic space that lie outside the atomic nucleus. Electrons are found in the energy levels. The energy levels are also called quantum levels.

  [i] Electrons

These subatomic particles carry a negative one (-1) electrical charge.

In a neutral atom of any element the number of protons equals the number of electrons.

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Atomic number:- identifies the element- equal to the number of protons in the nucleus- since atoms are electrically neutral, # of protons = # of electrons

Mass number:- # of protons + # of neutrons- # neutrons = atomic mass (rounded off) – atomic #- Protons and neutrons account for most of the mass of the atom

Electron Energy Diagrams for Atoms (Bohr diagrams): an energy level represents a specific value of energy of an electron and

corresponds to a general location around the nucleus the number of occupied energy levels in any atom is normally the same as

the period number in which an atom appears for the first 3 energy levels, the maximum number of electrons that can be

present are 2, 8 and8 in order of increasing energy (increasing distance from the nucleus)

a lower energy level is filled with electrons to its maximum level before the next level is started.

The electrons in the highest (outermost) occupied energy level are called valence electrons. Number of valence electrons is the same as the group number for group A elements (1,2,13-18).

Stable Atoms

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The outermost energy level of any atom is called the valence level. The electrons found in this level are called valence electrons. The number of valence electrons an atom has and the way they are arranged determines whether or not it will form a compound and the type of compound itwill become a part of.

have low chemical reactivity include noble gases, all of which have 8 valence electrons (except He,

which has 2) stability is a function of having a full complement of valence electrons. Atoms that do not have full electrons energy levels are

unstable and must gain, lose or share electrons to become stable. other atoms can become more stable by reacting and changing the number

of their electrons; thereby attaining the same stable electron configuration (structure) of the nearest noble gas:

atoms can follow one of 2 rules:a) Octet rule - atoms attempt to obtain 8 valence electrons

- includes most atomsb) Duet rule - atoms attempt to obtain 2 valence electrons

- includes H, Li and Be

Atoms can achieve a stable octet or duet by forming ions.

Ions

- an atom or groups of atoms that have a positive or negative charge, due to the loss or gain of one or more electrons.

- Single atoms form simple ions (monatomic ions); groups of atoms form complex ions (polyatomic ions)

Example: sodium metal and chlorine gas react to produce NaCl, a very stable and unreactive substance, compared to Na or Cl. The sodium atom loses 1 electron to the chlorine atom so both of their outer levels are filled. In doing so, the atoms form ions of opposite charge.

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How ionic bonds form:

Begin with atoms of two different elements that do not have 8 electrons in their outer most energy level.

The sodium atom (Na ) donates 1 electron becoming a positively sodium ion ( Na+).

The chlorine atom (Cl) accepts the donated electron becoming a negatively charged chloride ion (Cl-).

 

This chemical reaction produces table salt (NaCl).

Two types of ions:a) Cations:

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Ions form when atoms lose electrons Metals form cations Have a positive charge because they have more protons than electrons Name stays the same (sodium atom written as sodium ion)

b) Anions:

Ions form when atoms gain electrons Nonmetals form anions Have a negative charge because they have more electrons than protons Name changes to “-ide”. (chlorine atom becomes chloride ion)

Note:

1. Both cations and anions are more stable than the atoms from which they form since these ions have the same stable electron structure/configuration as the nearest noble gas.

2. Boron, carbon and silicon do not tend to form ions (they share electrons with other atoms)

3. The noble gases do not form ions since they are already stable (have filled orbitals)

4. Hydrogen can form a cation or an anion: Cation: H+, hydrogen ion has 1 proton but no electrons Anion: H-, hydride ion has 1 proton and 2 electrons

Types of Compounds

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1. Ionic Compounds:- involve the transfer of electron(s) between 2 oppositely charged ions

(cation and anion)- metal and a nonmetal or a combination involving a complex ion- forms an ionic bond - exists as an ionic crystal lattice (not individual mlecules

- known as a formula unit (eg. A formula unit of salt, not a molecule)

Formula Unit: a chemical formula showing the simplest whole number ratio of cations to anions in an ionic compound.

2. Molecular compounds

- involve the sharing of electrons between nonmetals- forms a covalent bond- exists as individual molecules

Properties of Ionic and Molecular Compounds

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1. State at room temperature:- all ionic compounds are solids- molecular compounds may be a solid, liquid or a gas

2. Conductivity of solution:- ionic compounds conduct electricity (they are electrolytes)- molecular compounds do not conduct electricity (non-electrolytes)

3. Solubility in water:- ionic compounds are soluble, to varying degrees (some better than others)

and form colored or colorless solutions.- molecular compounds may or may not be soluble

Nomenclature

Chemical nomenclature is the systematic naming of chemical compounds.

Compounds can be divided into two basic categories, those which are true binary compounds (they contain only two types of elements), and those which contain more than two different types of elements.

Rules for Naming ionic compounds:

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Identify the type of ions:

A. Monoatomic or simple ions

- Single atoms that have lost or gained one or more electrons- Form binary ionic compounds (2 simple ions)

- Eg. Sodium + chlorine Na+ Cl-

B. Polyatomic or complex ions

- cations or anions composed of a group of atoms with a net positive or negative charge.

- eg. Nitrate NO3-

Ammonium NH4+

C. Multivalent ions

- certain transition metals can form more than one type of ion, each with a different charge.

- the one written on top is the more common ion

- eg. Cu2+ or Cu+

D. Hydrates

- ionic compounds that contain water in their structure- eg. CuSO4 H2O

Naming Ionic compounds; (4 types)

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1. Binary Ionic compounds

Consist of cations and anions Cations are written first, anions are second (name changes to “-ide”) The total charge must be zero

Rules:i. write the symbols for the ions involved

eg. Silver and chlorine

Ag+ and Cl-

ii. determine the lowest whole number ratio of ions which will provide an overall net charge of zero

Ag1+ Cl1

- becomes AgCl (silver chloride)

iii. do not write charges in your final answer

iv. if one of your charges is odd, you can use the criss-cross method

Example:

potassium and oxygen: aluminum and sulfur

Potassium - K+ aluminum - Al3+ Oxygen - O2- sulfur - S2-

K2+ O1

- becomes K2O Al23+ S3

2- becomes Al2S3

Potassium oxide aluminum sulfide

Naming Ionic compounds: Name the cation (positive ion) by writing the full name of the metallic element Name the anion by abbreviating the nonmetallic element to “ide” (chlorine to

chloride.

Practice:

NaCl BaCl2 Al2O3

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2. Polyatomic/Complex Ions

Complex ions are groups of atoms that are made stable by sharing electrons and which then become even more stable by gaining (usually) or losing electrons. Unlike neutral molecules, complex ions carry an electric charge and do not exist by themselves. An ionic bond is formed by the attraction of a positive simple ion to a negative complex ion or of a positive complex ion (NH4

+) to a negative simple or complex ion. The total positive charge in the formula must be equal to the total negative charge.

Polyatomic ion: atoms of 2 or more elements covalently bonded together with an overall charge

eg. Nitrate NO3-

Ammonium NH4+

Rules: Don’t change the ending of a polyatomic ion! Name the cation, then name the anion Balance the charges If you need more than 1 complex ions, use brackets for that group

Example:

1. Sodium ions and carbonate ions bond ionically to form an ionic compound.Na+ CO3

2- Na2CO3 sodium carbonate

2. Ammonium ions and hydrogen phosphate ions bond ionically to form an ionic compound.

NH4+ HPO4

- (NH4)2HPO4 ammonium phosphate

3. magnesium ions and hydroxide ions bond ionically.

Mg 2+ OH- Mg(OH) 2

Practice

NaNO3

Al2(SO4)3

Mg(OH)2

NaCH3COO

3. Multivalent

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- Certain transition metals can form more than one type of ion, each with a different charge.

- The transition metals have various electron configurations that will make them stable.

- the one written on top is the more common ion.

eg. Cu2+ - copper (II)Cu+ - copper (I)

Fe3+ - iron (III)Fe2+ - iron (II)

- To name a multivalent metal compound, you must specify the charge on the ion as a roman numeral in brackets (Stock naming system).

- Eg. Iron (ii) oxide FeO Iron (iii) oxide Fe2O3

- Do not write charges as a roman numeral if the metal is not multivalent.

Practice:CuSO4

PbO

uranium (vi) oxide

uranium (iv) oxide

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4. Hydrates

- A number of ionic compounds called hydrates produce water when they decompose upon heating.

-eg. CuSO4 5H2O copper (II) sulfate pentahydrate

- When the formula of a hydrated compound is written, the number of water molecules is also included. For example, copper (II) sulfate pentahydrate is written as CuSO4 5 H2O, meaning 5 molecules of water are bonded within the ionic crystal for every one formula unit of CuSO4.

- Rules for naming:a) Name the ionic part of the formula firstb) Name the water part second, using a prefix system for the number of

water moleculesc) Add prefix to “hydrate”d) Prefixes: mono - 1

Di - 2Tri - 3Tetra - 4Penta - 5Hexa - 6Hepta - 7Octa - 8Nona - 9Deca - 10

Example:Barium hydroxide hexahydrate Ba(OH) 2 6 H2O

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MOLECULAR COMPOUNDS

Binary molecular compounds form between 2 non-metals Covalent bonds: shared electrons Molecular formula: shows number and kind of atoms in a molecule

Rules for naming molecular compounds: Write the name of the first element of the formula in full Second element ends with “-ide” Use prefixes to specify number of atoms of each element in the molecule; use the

same prefixes used for hydrates No charges used in formula The prefix “mono-” is optional on the first name

Name the following: NO nitrogen monoxide CO2 carbon dioxide N4O9 tetranitrogen nonaoxide N6O hexanitrogen monoxide

Write formulas for the following: Boron trifluoride Sulfur hexafluoride Nitrogen monoxide Phosphorous pentachloride

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The following trivial names for molecular compounds must be memorized:

Name Formula Name FormulaWater H2O Propane C3H8

Hydrogen peroxide

H2O2 Ammonia NH3

Glucose C6H12O6 Methane CH4

Sucrose C12H22O11 Methanol CH3OH

Ozone O3 Ethanol C2H5OH

Molecular Elements:

Some non-metals do not exist as single atoms in nature. Some are diatomic (containing 2 atoms) and some are polyatomic (more than 2 atoms):

Element Formula Element FormulaPhosphorus P4(s) Nitrogen N2(g)

Sulfur S8(s) Fluorine F2(g)

chlorine Cl2(g) Iodine I2(g)

Bromine Br2(l) oxygen O2(g)

hydrogen H2(g)

Note:P.S. Clem Brown Has No Friends In Ottawa.

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ACIDS

They must contain hydrogen (H+)

They must be dissolved in water (aqueous, aq); the formula will always contain the subscript aq.

Rules for naming acids:

Hydrogen is always the positive ion for an acid

Ending Acid Name 1. -ide begins with hydro, ends with -ic and acid

2. –ite ends with –ous and acid

3. –ate ends with –ic and acid

1. hydrogen _____ide becomes hydro____ic acidEx: hydrogen chloride becomes hydrochloric acid

HCl HCl(aq)

2. hydrogen _____ate becomes______ic acidEx: hydrogen sulfate becomes sulfuric acid

H2SO4 H2SO4(aq)

3. hydrogen _____ite becomes____ous acidEx: hydrogen nitrite becomes nitrous acid

HNO2 HNO2(aq)

NOTE: when sulf is the root we add “ur” and when phosph is the root weadd “or” to make it sound better.

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Chemical Reactions

A chemical reaction represents a chemical change.

A + B ® C + D

Reactants Products

Note:Reactants are your starting materialsThe arrow means producesProducts are the new substances formed

A chemical equation is a shorthand way of representing what experimental evidence indicates happens in a chemical reaction and must show the substances involved and the correct number of each atom or ion involved.

A chemical equation only involves a rearrangement of the atoms involved. It does not produce new atoms, only new substances. This is according to the Law of conservation of matter (or mass) – matter is neither created or destroyed in a chemical change.

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Example:

Paragraph form: iron reacts with oxygen to produce iron (iii) oxide.

Word equation:iron + oxygen ® iron (iii) oxide

Skeleton Chemical equation: Fe(s) + O2(g) ® Fe2O3(s)

Balanced chemical equation:4Fe(s) + 3O2(g) ® 2Fe2O3(s)

How to write equations:

Step 1: Write the correct formula and state (s, l, g or aq) for each reactant and product

Step 2: Balance the atoms without changing the formula of the compounds. Change the number in front of the formula, called the coefficient, to show the correct number of atoms involved. Do not show “1” in the balancing.

Generally, start by balancing the atom of which there is the greatest number. Find the lowest common multiple of the numbers of reactants and product atoms. Then balance the rest of the atoms. Leave the elements until the end.

Example: Write the equation for magnesium reacting with oxygen. (Reminder of P.S.Clem Brown Has No friends in Ottawa).

Magnesium + oxygen ® magnesium oxide

Mg(s) + O2(g) ® MgO(s)

2Mg(s) + O2(g) ® 2MgO(s)

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Types of Reactions:

1. Simple composition/formation

- 2 or more elements combine to form a compound- Element 1 + element 2 ® Compound- Eg. Aluminum + oxygen ® aluminum oxide

4 Al(s) + 3 O2(g) ® 2 Al2O3(s)

2. Simple Decomposition

- compound ® element + element- eg. Sodium chloride ® sodium + chlorine

2 NaCl(s) ® 2 Na(s) + __Cl2(g)

3. Single replacement

- element 1 + compound 1 ® element 2 + compound 2- eg. Magnesium reacts with hydrochloric acid.

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

- Use the solubility chart to predict the states of the products if at least one of your reactants is aq.

4. Double replacement

- Compound 1 + compound 2 ® compound 3 + compound 4- Eg. Potassium iodide reacts with lead (ii) nitrate).

KI(aq) + Pb(NO3)2(aq) ® PbI2(s) + KNO3(aq)

A special type of double replacement is a neutralization reaction, where an acid joins with a base to produce water and a “salt”

Acid + base ® water + “salt”

5. Combustion

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Hydrocarbon + oxygen ® carbon dioxide + water

Eg. Burning propane

C3H8(g) + O2(g) ® CO2(g) + H20(g)

6. Other

The reaction is classified as “other” if it is not one of the above.

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