An Introduction to Molecular Elements and Compounds.pdf

Molecular
Elements and
Compounds

Molecular Elements
A molecular element consists of
molecules containing atoms of
only one type of element.
Nitrogen, N2(g), and bromine,
Br2(l), are diatomic elements:
each molecule is made up of 2
atoms. Ozone, O3(g), and
sulfur, S8(s), are other
examples of molecular
elements. All the molecules in a
particular molecular element are
identical.

Molecular Compounds
A molecular compound is also made up of
molecules, each with the same arrangement of
specific atoms. Each molecule in a molecular
compound is made up of at least 2 different
elements—sometimes many more. All the
molecules in a particular compound are identical,
but different from the molecules in another
compound.
Molecular compounds generally consist of only
non-metallic elements, although some important
molecular compounds involve metals and
metalloids.

Properties of Molecular Substances
▪ Molecular substances exist as gases, liquids, and solids at ambient
temperatures
▪ They have low melting point due the weak intermolecular forces which
break easily.
▪ They are non-conductors of heat and electricity because of the absence
of free electrons or ions.
▪ They are soluble in organic solvents but not in water.

Formation of Molecular Compounds
▪ When two or more non-metals combine, they share one or more pairs of electrons.
▪ A shared pair of electrons is called a covalent bond.
▪ The share electrons in the covalent bond usually arise from the outer shell of the atoms
that combine.
▪ When some non-metal atoms combine, not all the electrons in the outer shell form
covalent bonds. The pairs of electrons not used in covalent bonding are called lone
pairs.
▪ In writing displayed formulae, a single covalent bond is shown as a single line between
the atoms, e.g. Cl – Cl.

Formation of Molecular Elements
When two elements that have a similar ability to attract electrons collide, they will both attempt to pull in the
electrons of the other element. This will result in a tug of war situation however neither element will be able to
completely pull the other electrons into its structure, the electrons are therefore shared since neither element is
able to fully pull in electrons.
The sharing of electrons allows each atom to achieve a full valence shell. This full valence shell is a stable
electron arrangement similar to that of the nearest noble gas. The nearest noble gas to hydrogen on the

Bonding Electrons
▪ Atoms typically have unpaired electrons that are available to form a covalent bond with
another atom. These electrons are called bonding electrons.
▪ The number of bonding electrons in an atom affects the number of bonds it can form.
This number is known as the bonding capacity.

Lewis Structure for Covalent
Compounds
Lewis's structure is a representation of covalent bonding based on Lewis symbols; a
model in which shared electron pairs are shown as lines and unshared electrons are
shown as dots.
A single dash represents a single covalent bond. A double dash represents a double
covalent bond. A double covalent bond is formed when two atoms share two pairs
of electrons. A triple dash represents a triple covalent bond, which is formed when
two atoms share three airs of electrons.

An Introduction to Molecular Elements and Compounds.pdf

Molecular Compounds
Molecular compounds form in much the same way as molecular elements. The difference
is that the atoms are not from the same element.

Molecular
Compounds
Molecules are not limited to two
atoms each. Example CO2 and H2O

Lewis Structure
The Lewis structure for molecular
compounds requires both the lone
pairs of electrons surrounding each
atom and the covalent bonds (single,
double, or triple) between the atoms.
The Lewis structure can be simplified
to a structural formula diagram that
only includes the covalent bonds in the
diagrams.

Steps in Constructing Lewis Structure
1. Arrange the symbols with the element having the largest bonding capacity in the center, spreading
surrounding atoms as far apart as possible.
2. Count and record the total number of valence electrons available.
3. Place single covalent bonds (2 electrons) between the central atom and surrounding atoms.
4. Distribute remaining electrons as lone pairs to satisfy the octet rule for outer atoms (except
hydrogen).
5. Subtract used electrons from the total and record the remaining number.
6. Place any remaining electrons on the central atom in pairs.
7. If the central atom lacks a full octet, move lone pairs from surrounding atoms to form double or triple
bonds.
8. Verify that all atoms (except hydrogen) have complete octets.
9. Replace shared electron pairs with dashes to represent covalent bonds, using double or triple
dashes as needed.

Lewis Structure
Draw a Lewis structure and structural formula for each of the following molecules:
(a) F2
(f) C2
H6
(b) N2
(c) PH3 (h) CO (d) CO2

Lewis’s Structure and Structural Formula of the
Following Polyatomic ions:
(a) OH−
(Hydroxide)
(b) CN−
(Cyanide)
(c) NH4+
(Ammonium)
(d) PO43−
(Phosphate)

Electronegativity:
Electronegativity refers to an atom’s ability to attract bonding electrons towards itself when
forming a chemical bond. This concept was introduced by Linus Pauling, who also
developed a numerical scale to compare the electronegativities of elements.
Fluorine (F) has the highest electronegativity (4.0), meaning it has the strongest ability to
attract electrons.
Francium (Fr) has the lowest electronegativity (0.7), meaning it has the weakest ability to
attract electrons.

Electronegativity:
Electronegativity increases across a period (left to right) and decreases down a group (top
to bottom) in the periodic table.
This trend is closely related to atomic radius:As the nucleus becomes increasingly
shielded by inner electron shells, the power of the nucleus to draw electrons to itself
lessens.
Smaller atoms (like fluorine) have a stronger pull on shared electrons.Larger atoms (like
francium) have a weaker attraction because their valence electrons are farther from the
nucleus.
Since electronegativity cannot be measured directly, scientists calculate it based on other
atomic properties, such as ionization energy.

TRENDS IN GROUPS
ELECTRONEGATIVITY
DECREASES

How Electronegativity Determines Bond
Type
The electronegativity difference (ΔEN) between two atoms determines the nature of their bond.
It helps predict whether a bond is ionic, polar covalent, or non-polar covalent.
1. Ionic Bonds (ΔEN ≥ 1.7)
Occurs when the electronegativity difference is large (e.g., NaCl: ΔEN = 2.3).The more
electronegative atom completely takes an electron from the other atom, forming ions.
Example: Sodium chloride (NaCl)
Na (0.9) loses an electron → Na⁺
Cl (3.2) gains an electron → Cl⁻

Non-Polar Covalent Bonds (ΔEN = 0)
Occurs when atoms have identical (or nearly identical) electronegativity.Electrons are
shared equally.
Example: Cl₂ (Chlorine gas)
Both Cl atoms have the same electronegativity (ΔEN = 0).The bond is non-polar, meaning
no partial charges develop.

Polar Covalent Bonds (0 < ΔEN < 1.7)
Occurs when one atom is slightly more electronegative than the other.Electrons are
unequally shared, creating a dipole (partial positive and negative charges).
Example: HCl (Hydrogen chloride)
Cl (3.2) is more electronegative than H (2.2).
Cl pulls electrons closer, creating a partial negative charge (δ⁻), while H has a partial
positive charge (δ⁺).This results in a polar bond.

Quick Check
Explain how the trends of electronegativity in the periodic table relate to those of atomic radius.
Using only their relative position on the periodic table, arrange the following elements in order of
increasing electronegativity: K, Cs, Br, Fe, Ca, F, Cl0.
Describe how we can use electronegativity values to predict the types of bonds that will form in a
compound.
Distinguish between a polar covalent bond and a non-polar covalent bond.
For each bond listed, determine the electronegativity difference and predict what type of bond
(non-polar covalent, polar covalent, or ionic) would form between the two elements:
(a) Ca–S (b) H–F (c) P–H (d) C–Cl (e) C–O (f) Li–Cl

Electronegativity
•Ca (Calcium) = 1.00
•S (Sulfur) = 2.58
•H (Hydrogen) = 2.20
•F (Fluorine) = 3.98
•P (Phosphorus) = 2.19
•C (Carbon) = 2.55
•Cl (Chlorine) = 3.16
•O (Oxygen) = 3.44
•Li (Lithium) = 0.98

Chemical Reactions
The change of one or more substances into other substances that have different
composition and properties is called a chemical reaction.
Signs that a chemical reaction is occurring may include:
▪ a colour change
▪ a change in temperature
▪ effervescence ( a gas is given off)
▪ • A precipitate forms.
▪ A chemical equation tells us what chemical changes take place during a reaction. It tells
us what the reactants (things that react) and what the products (things that are formed)
are.
▪ A chemical equation must be balanced. This means that the total number and types of
atoms on the right side of the equation must be equal to those on the left side of the
equation. This is because atoms cannot be created or destroyed.

The Law of Conservation of Matter in
Equations
The law of conservation of matter states that in a chemical reaction, the mass of the
products is equal to the mass of the reactants.
In a chemical reaction, some of the bonds in the reactants break and new bonds are made
in forming the products. The atoms or ions rearrange themselves so that there is the same
number of each type of atom on each side of the equation.
2HCl(aq) + CaCO3
(s) 🡪 CaCl2
(aq) + H2
O(l) + CO2
(g)

State Symbols
The state symbols tell us the physical states of the reactants and products in a chemical
reaction.
(s) 🡪 solid state
(l) 🡪 liquid state
(g) 🡪 gaseous state
(aq) 🡪 aqueous state (solution in water)

Synthesis Reactions
In a synthesis reaction, two reactants combine to form one larger or more complex
product. For this to occur, the reactants must first collide, break existing bonds between
their atoms, and form new bonds. The chemical equations for synthesis reactions fit the
general pattern:
A + B 🡪 AB
2 H2(g) + 0 2(g) 🡪 2 H20(g)
A reaction in which two or more substances (elements or compounds) combine
together to form a new substance.

Predicting the Product of Synthesis
Reactions
Step 1. Identify the reactants.
Step 2. Identify the type of reaction.
Step 3. Use the ionic charges of the reactants to predict the formula of the product. Reduce to simplest ratios, if
necessary.
Step 4. Balance the chemical equation and write the balanced chemical equation, including state symbols.
▪ Practice
▪ Write the balanced chemical equation for the reaction of these pairs of reactants:
▪ (a) calcium and bromine
▪ (b) aluminum and oxygen

Criss Cross the Charges to Balance
▪ Mg + N 🡪 ?
▪ Mg = +2
▪ N = -3
▪ Mg N
+2 -3
▪
▪
▪ Synthesis: A + B = AB
Mg3N2

Synthesis Reactions of Non-Metals
The products for the synthesis reactions of nonmetals will always be a molecular
compound.
Synthesis Reactions with Hydrogen
Hydrogen is different from other elements in its group because it usually forms molecular
compounds (made of atoms) instead of ionic compounds (made of charged
particles).Even though hydrogen doesn’t always form ions, we can still use ionic charges
to predict what happens in chemical reactions.
For example: Hydrogen and Chlorine Reaction or Hydrogen and Oxygen

Reactions Involving Non-Metals (Other
than Hydrogen)
Predicting the products of synthesis reactions involving non-metals can be challenging, as
the outcome often depends on reaction conditions.
For instance, when carbon reacts with oxygen:
In an oxygen-rich environment, carbon dioxide (CO₂) is formed:
C + O₂ → CO₂
Under oxygen-limited conditions, carbon monoxide (CO) is produced:
2 C + O₂ → 2 CO
The exact product can only be determined through experimental analysis.

Synthesis Reactions Involving
Compounds
▪ Synthesis reactions are not limited to elemental reactants—compounds can also undergo synthesis to form new
substances.
Carbon Dioxide and Water: Formation of Carbonic Acid
• When carbon dioxide (CO₂) dissolves in water, a chemical reaction occurs, producing carbonic acid (H₂CO₃):
H₂O + CO₂ → H₂CO₃
• The presence of carbonic acid can be confirmed using bromothymol blue, an acid-base indicator that shifts from
blue to yellow in acidic conditions.
▪ Sulfur Trioxide and Water: Formation of Sulfuric Acid
Sulfur trioxide (SO₃) is generated during fossil fuel combustion.
Upon reacting with water, it forms sulfuric acid (H₂SO₄):
H₂O + SO₃ → H₂SO₄

Examples of Different Types of Synthesis
Reactions
1. Two Compounds → New Compound
•Example: Water + Carbon Dioxide → Carbonic Acid
H₂O + CO₂ → H₂CO₃
2. Metal Oxide + Water → Metal Hydroxide
•Example: Calcium Oxide + Water → Calcium Hydroxide
CaO + H₂O → Ca(OH)₂
3. Non-Metal Oxide + Water → Acid
•Example: Sulfur Trioxide + Water → Sulfuric Acid
SO₃ + H₂O → H₂SO₄
4. Metal Oxide + Non-Metal Oxide → Salt
•Example: Calcium Oxide + Carbon Dioxide → Calcium Carbonate
CaO + CO₂ → CaCO₃
•Acid + Base → Salt + Water
5. Example: Hydrochloric Acid + Sodium Hydroxide → Sodium Chloride + Water
HCl + NaOH → NaCl + H₂O

Decomposition Reactions
A decomposition reaction is a type of chemical reaction in which a single compound breaks down into two or
more simpler substances. This is the opposite of a synthesis reaction, where simple reactants combine to form
a more complex product.
General Pattern: AB → A + B
In this pattern, a compound (AB) decomposes to form two or more products (A and B).
Important Points
Energy Requirement:
Decomposition reactions often require energy to initiate, even if the reaction releases energy afterward. For
example, in the case of TNT (trinitrotoluene), a small electric current is used to trigger the reaction.
Product States:
Decomposition reactions can produce various products, including gases, solids, or liquids, depending on the
nature of the compound and the conditions of the reaction.Many decomposition reactions of explosive
compounds release gaseous products. These gases expand rapidly when heated, creating a powerful
destructive force.

General Guidelines for Predicting
Decomposition Reactions
1. Decomposition of Ionic Compounds:
1. Simple ionic compounds, like potassium chloride (KCl), can often be predicted to decompose into their elements when heated or
subjected to an electrical current.
1. Example:
2 KCl(l) → 2 K(s) + Cl₂(g)
2. Decomposition of Metal Carbonates:
1. Metal carbonates (e.g., calcium carbonate) often decompose into metal oxides and carbon dioxide (CO₂) gas when heated.
1. Example:
CaCO₃(s) → CaO(s) + CO₂(g)
3. Decomposition of Metal Hydroxides:
1. Metal hydroxides (e.g., sodium hydroxide) usually decompose into metal oxides and water when heated.
1. Example:
2 NaOH(s) → Na₂O(s) + H₂O(g)
4. Decomposition of Metal Chlorates:
1. Metal chlorates, like potassium chlorate (KClO₃), decompose to form metal chlorides and oxygen gas (O₂) when heated.
1. Example:
2 KClO₃(s) → 2 KCl(s) + 3 O₂(g)
5. Decomposition of Molecular Compounds:
1. Molecular compounds (especially complex ones) are harder to predict, as their decomposition often depends on the conditions of the
reaction.

1. Classify these reactions as Synthesis or Decomposition.
▪ Justify your choice.
• (a) 2 Al + 3 Br₂ → 2 AlBr₃
• (b) 2 HCl → H₂ + Cl₂
• (c) CaO + H₂O → Ca(OH)₂
• (d) P₄ + 5 O₂ → 2 P₂O₅
2. Predict the products of these synthesis or decomposition reactions. Write a balanced chemical equation to represent each reaction.
• (a) Z + S → ?
• (b) CaCl₂ → ?
• (c) NH₃ + HCl → ?
• (d) K₂O → ?
• (e) AlCl₃ → ?
• (f) Mg(OH)₂ → ?
3. Hydrogen peroxide forms gas bubbles when it is added to blood. The other reaction product is water. Inserting a glowing splint into a sample of this gas causes the splint to
relight.
• (a) Identify the gas.
• (b) Classify the reaction.
• (c) Write a balanced chemical equation for this reaction.
4. Write a balanced chemical equation for each of these reactions:
• (a) Aluminum metal readily reacts in air to form a hard protective coating of aluminum oxide.
• (b) Copper(II) oxide and carbon dioxide are produced when copper(II) carbonate is heated.
• (c) Solid nitrogen triiodide is a shock-sensitive explosive that is stable when wet and explosive when dry. This compound decomposes rapidly to produce a gas when detonated.

Single Displacement Reactions
• A single displacement reaction is a type of chemical reaction in which one element
displaces or replaces another element in a compound, producing a new compound and
a new element.
• General pattern:
A + BC → AC + B
Where:
• A is an element.
• BC is a compound.
• AC is the new compound.
• B is the displaced element.

Single Displacement Reactions
In this type of reaction, one element replaces another element in a compound. This results
in:
• A new compound.
• A displaced element that becomes free.
▪ Example:
• The reaction of zinc (Zn) with copper(II) sulfate (CuSO₄) displaces copper (Cu) to form
zinc sulfate (ZnSO₄) and copper metal (Cu).
• Balanced equation:
Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

Activity series
•This is a list of metals ranked by their reactivity in single displacement reactions.
•The more reactive a metal, the higher it is on the activity series.
•High-reactivity metals: Lithium (Li), Potassium (K), Calcium (Ca).
•Low-reactivity metals: Gold (Au), Silver (Ag), Copper (Cu).
•Important rules:
1. A more reactive element can displace a less reactive element from a compound, but a
less reactive element cannot displace a more reactive one.
2. The farther apart two elements are on the activity series, the faster the reaction will
likely occur

Using the Activity Series to Predict
Reactions:
• To predict if a reaction will occur, check if the element is higher in the activity series than
the metal or hydrogen it’s displacing.
• Example: Magnesium (Mg) is more reactive than copper (Cu), so it will replace copper in
copper(II) sulfate.
Mg(s) + CuSO₄(aq) → MgSO₄(aq) + Cu(s)
• If the element is lower than the metal in the compound, no reaction occurs.
• Example: Lead (Pb) is less reactive than zinc (Zn), so no reaction will occur when lead is added to
zinc nitrate.

Reactions Involving Metals and Water or
Acids:
• Metals also react with acids or water in single displacement reactions.
• Metals react with acids:
• Example: Magnesium reacts with hydrochloric acid to produce magnesium chloride and hydrogen
gas.
Mg(s) + 2 HCl(aq) → MgCl₂(aq) + H₂(g)
• Metals react with water:
• Example: Sodium reacts with water to form sodium hydroxide and hydrogen gas.
2 Na(s) + 2 H₂O(l) → 2 NaOH(aq) + H₂(g)

The Halogen Activity Series:
• Halogens (Group 17 elements) also have an activity series based on their reactivity in
displacement reactions.
• More reactive halogens can displace less reactive halogens from compounds:
• Fluorine is the most reactive, followed by chlorine, bromine, and iodine.
• Example: Chlorine gas displaces iodine from potassium iodide solution:
Cl₂(g) + 2 KI(aq) → 2 KCl(aq) + I₂(aq)
• No reaction occurs when iodine is added to potassium chloride because iodine is less
reactive than chlorine.

Use the activity series to predict whether the
following reactions will occur:
1. Magnesium (Mg) + Silver Nitrate (AgNO₃)
2. Zinc (Zn) + Iron(II) Chloride (FeCl₂)
3. Nickel (Ni) + Aluminum Nitrate (Al(NO₃)₂)
4. Sodium (Na) + Water (H₂O)

•Use the activity series of metals to predict whether each of these reactions
occurs. Write a balanced chemical equation for each reaction that does
occur. Write no reaction for those that do not.
(a) AI(s) + AgNO₃(aq) →
(b) Zn(s) + Pb(NO₃)₂(aq) →
(c) Au(s) + H₂O(l) →
(d) Mg(s) + H₂SO₄(aq) →
(e) Ca(s) + H₂O(l) →
(f) Al(s) + HCl(aq) →
•Use the halogen activity series to predict whether each of these reactions
occurs. Write a balanced chemical equation for each reaction that does
occur. Write no reaction for those that do not.
(a) Br₂(l) + NaI(aq) →
(b) Cl₂(g) + KF(aq) →
(c) F₂(g) + CaBr₂(aq) →

Displacement Reactions
A double displacement reaction is a chemical reaction where
elements from two compounds switch places, forming two
new compounds.
General equation:
𝐴𝐵+𝐶𝐷→𝐴𝐷+𝐶𝐵
Example reaction:𝐶𝑢𝑆𝑂4(𝑎𝑞)+2𝑁𝑎𝑂𝐻(𝑎𝑞)→𝐶𝑢
(𝑂𝐻)2(𝑠)+𝑁𝑎2𝑆𝑂4(𝑎𝑞)
In this reaction, copper and sodium ions exchange places,
forming copper(II) hydroxide precipitate.

Precipitation Reactions
Precipitation Reactions occur when a reaction forms an insoluble product (precipitate).
Precipitate: An insoluble solid that forms in a solution during a chemical reaction.
Solubility: The ability of a substance to dissolve in a solvent.
Key concepts:
Solute: The substance that dissolves.
Solvent: The substance that dissolves the solute. Very soluble substances remain in
solution, while slightly soluble substances form precipitates.

Sample Reaction
▪ When potassium chloride and silver nitrate are combined:
• Reactants: KCl (aq) + AgNO₃ (aq)
• Ion Separation: K⁺, Cl⁻, Ag⁺, NO₃⁻
• Possible Products: KNO₃, AgCl
• Solubility Check: AgCl is slightly soluble (precipitate), while KNO₃ is very soluble.
• Balanced Equation:
• KCl(aq)+AgNO3(aq)→AgCl(s)+KNO3(aq)

Sample Reaction Polyatomic Ions
Determine if a precipitate will form when solutions of potassium sulfate and iron(III) chloride are combined
Step 1. Write the chemical formulas of the reactants.
▪ Step 2. Separate the reactants into their ions.
K₂SO₄ → 2K⁺ + SO₄²⁻
FeCl₃ → Fe³⁺ + 3Cl⁻
▪ Step 3. Combine cations and anions to form new compounds.
Possible products:
• KCl (from K⁺ and Cl⁻)
• Fe₂(SO₄)₃ (from Fe³⁺ and SO₄²⁻)
▪ Step 4. Check for solubility.
Both KCl and Fe₂(SO₄)₃ are soluble, so no precipitate forms.

Check if Precipitates are Formed
(a) Na₂S(aq) + Pb(NO₃)₂(aq) → ?
(b) NH₄Cl(aq) + K₂SO₄(aq) → ?
(c) FeCl₃(aq) + Na₂CO₃(aq) → ?

Gas-Producing Reactions
• Occur when double displacement reactions generate a gas.
• Some reactions produce gas directly, while others form an unstable product that decomposes into a gas.
• Examples:
• Direct Gas Formation:
• K2S(aq)+2HCl(aq)→H2S(g)+2KCl(aq)
• Produces hydrogen sulfide gas (H₂S).
• Gas from Unstable Products:
• MgCO3(aq)+H2SO4(aq)→MgSO4(aq)+H2CO3(aq)
• H₂CO₃ decomposes into CO₂ gas and water:
• H2CO3(aq)→H2O(l)+CO2(g)
• Overall Reaction:
• MgCO3(aq)+H2SO4(aq)→MgSO4(aq)+H2O(l)+CO2(g)

Neutralization Reactions
Occur when an acid reacts with a base, producing water and a salt.
Example:𝐻𝑁𝑂3(𝑎𝑞)+𝑁𝑎𝑂𝐻(𝑎𝑞)→𝐻2𝑂(𝑙)+𝑁𝑎𝑁𝑂3(𝑎𝑞)
Antacid Reaction
Example: Magnesium hydroxide neutralizes stomach acid:
𝑀𝑔(𝑂𝐻)2(𝑠)+2𝐻𝐶𝑙(𝑎𝑞)→2𝐻2𝑂(𝑙)+𝑀𝑔𝐶𝑙2(𝑎𝑞)

Salts Explored
Salts are ionic compounds that result from the reaction between an acid and a base during neutralization. In this
reaction, the acid donates H⁺ (hydrogen ions), and the base donates OH⁻ (hydroxide ions). These combine to
form water (H₂O), while the remaining ions from the acid and base combine to form a salt.
Formation of Salts in Neutralization Reactions
▪ A neutralization reaction follows this general equation:
Acid+Base→Salt+Water
For example:
Hydrochloric acid + Sodium hydroxide → Sodium chloride + Water
HCl+NaOH→NaCl+H2OHCl + NaOH
Sulfuric acid + Magnesium hydroxide → Magnesium sulfate + Water
H2SO4+Mg(OH)2→MgSO4+2H2O
In each case, the cation (positive ion) comes from the base, and the anion (negative ion) comes from the acid,
forming the salt.

▪ How do single and double displacement reactions differ?
▪ Classify the reactions represented by these equations as either single or double
displacement:
HI(aq) + AgNO₃(aq) → AgI(s) + HNO₃(aq)
Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
ZnS(s) + 2 HCl(aq) → ZnCl₂(aq) + H₂S(g)
Cl₂(g) + 2 NH₄Br(aq) → Br₂(l) + 2 NH₄Cl(aq)
▪ Write a chemical equation for the reaction of barium carbonate with sulfuric acid:
▪ Write the chemical formula for each of the following compounds and predict their solubility in
water:
(a) Lead sulfate (b) Ammonium phosphate (c) Calcium sulfate
▪ Complete the chemical equations of the following reactions and indicate the state of each
compound:
ZnCl₂(aq) + 2KOH(aq) →
Ni(NO₃)₂(aq) + Na₂CO₃(aq) →
Ba(OH)₂(aq) + K₂SO₄(aq) →
FeSO₄(aq) + K₃PO₄(aq) →
▪ Will a precipitate form when the following solutions are combined?
BaCl₂(aq) + Ca(OH)₂(aq) →

Key Terms
▪ Combustion: Reaction where a substance reacts with oxygen, producing heat, light, and
new chemical compounds. It’s often exothermic (releases energy).
▪ Greenhouse Gas: Gases that trap heat in Earth’s atmosphere. Common ones include
CO₂, CH₄, H₂O vapor. Excess greenhouse gases contribute to global warming.
▪ Organic Compound: Molecules with carbon-carbon (C–C) or carbon-hydrogen (C–H)
bonds. Hydrocarbons are a subset of organic compounds.
▪ Air Pollution: Presence of harmful substances in the air, such as CO, NOₓ, SO₂, and
particulates, that can harm health or the environment.

Hydrocarbons
Molecules made up of carbon and hydrogen only.
They’re a primary component of fossil fuels: natural gas, gasoline, propane, etc.
Examples:
Methane (CH₄): Found in natural gas.
Propane (C₃H₈): BBQ fuel.
Butane (C₄H₁₀): Lighter fluid.
Ethyne (C₂H₂): Used in welding torches.
Heptane (C₇H₁₆): Component of gasoline.
Hexane (C₆H₁₄): Solvent and fuel.

Complete Combustion
Occurs with plentiful oxygen.
Produces carbon dioxide (CO₂), water (H₂O), and energy.
Releases maximum energy; burns cleanly with a blue flame.
Example (methane):
CH₄ + 2 O₂ → CO₂ + 2 H₂O + energy
Example (ethyne):
2 C₂H₂ + 5 O₂ → 4 CO₂ + 2 H₂O + energy

Incomplete Combustion
Happens with limited oxygen (fuel-rich conditions).
Products may include carbon monoxide (CO), soot (C), carbon dioxide (CO₂), water (H₂O), and less
energy.
Produces cooler, sooty, yellow flames.
Example (heptane):
C₇H₁₆ + 7 O₂ → 3 C + 2 CO + 2 CO₂ + 8 H₂O + energy
2 C₇H₁₆ + 11 O₂ → 14 CO + 8 H₂O + energy

Hazards of Incomplete Combustion
▪ Carbon monoxide: Binds to hemoglobin in red blood cells, preventing
oxygen transport—can be fatal in enclosed spaces.
▪ Soot: Microscopic carbon particles can damage lungs and reduce air
quality.
▪ Energy Waste: Less energy is produced compared to complete
combustion.

Quick Check
Write the chemical equation for the complete combustion of:
● a) propane
● b) butane
● c) octane
Write the chemical equation for the incomplete combustion of:
● a) propane (producing carbon monoxide)
● b) butane (producing carbon monoxide and carbon)
Name two products of complete combustion and explain why they are not pollutants.
Name three possible products of incomplete combustion and explain the dangers of each.

Oxides
● Definition: Oxides are binary compounds that contain oxygen and one other element.
● They form through direct reactions with oxygen, often from the air.
○ Fast reactions (e.g. burning magnesium)
○ Slow reactions (e.g. metal tarnishing or rusting)
Example Reaction: Combustion of Magnesium
2Mg(s)+O2(g)→2MgO(s)
● This reaction releases a large amount of energy.
● Produces magnesium oxide, a white powder.

Acid-Base
Acids:
● Release H⁺ (hydrogen ions) in water.
● Example:
HCl(aq)→H+(aq) + Cl−(aq)
● These H⁺ ions cause the solution to behave acidically.
Bases:
● Release OH⁻ (hydroxide ions) in water.
● Example:
NaOH(s)→Na+(aq)+OH−(aq)
pH Scale:
● Ranges from 0 (strongly acidic) to 14 (strongly basic).
● 7 = Neutral (e.g., pure water).
● The more H⁺ → the lower the pH → more acidic.
● The more OH⁻ → the higher the pH → more basic.

Oxides: Acidic vs. Basic
Non-Metal Oxides = Acidic Oxides
● Generally covalent and form acidic solutions in water.
● They lower pH when dissolved.
● Examples:
○ CO₂ forms carbonic acid:
CO2(g)+H2O(l)→H2CO3(aq)
○ NO₂ forms nitric and nitrous acids:
2NO2(g)+H2O(l)→HNO3(aq)+HNO2(aq)
○ SO₃ forms sulfuric acid:
SO3(g)+H2O(l)→H2SO4(aq)

Environmental Effects of Acidic Oxides
● These acids cause acid rain (pH < 5.6).
● Damage to:
○ Aquatic life (acidic lakes and streams)
○ Soil chemistry (leaching nutrients)
○ Man-made structures (corroding buildings, statues)

Metallic Oxides
Metal Oxides = Basic Oxides
● Generally ionic and form alkaline (basic) solutions in water.
● They increase pH when dissolved.
● Examples:
○ Sodium oxide (Na₂O):
Na2O+H2O→2NaOH(aq)
○ Calcium oxide (CaO):
CaO+H2O→Ca(OH)2(aq)

Industrial & Environmental Uses:
● Used to neutralize acid spills and acidic lakes.
● Lime (CaO) is used:
○ In cement, glass, soil treatment, and neutralization of acidic conditions.
○ By hazmat teams and farmers.

Neutralization
Neutralization reactions occur when an acid reacts with a base, resulting in a solution with a pH
closer to 7 than its reactants. Most follow this general formula:
Acid + Base → Water + Ionic Compound (Salt)
Some also produce carbon dioxide if a carbonate is involved:
Acid + Carbonate → Water + CO₂ + Ionic Compound

Common Neutralization Reactions
1. With Hydroxide Compounds
Example:
HCl(aq) + NaOH(aq) → H₂O(l) + NaCl(aq)
● Acid and base neutralize to form water and salt.
2. With Carbonate Compounds
Example:
H₂SO₄(aq) + CaCO₃(s) → H₂O(l) + CO₂(g) + CaSO₄(aq)
● Useful in nature and industry (e.g., coral reefs, antacids, lake rehabilitation).

Neutralization in Daily Life
1. Antacids and Heartburn Relief
● Stomach acid: Mainly HCl with pH ~1.5.
● Antacid ingredients:
○ Sodium bicarbonate (NaHCO₃)
○ Calcium carbonate (CaCO₃)
Examples:
● HCl + NaHCO₃ → NaCl + H₂O + CO₂
● 2 HCl + CaCO₃ → CaCl₂ + H₂O + CO₂
2. Lake Rehabilitation (Liming Lakes)
● Acid rain from SOₓ and NOₓ emissions lowers lake pH.
● Solution: Add calcium oxide (lime) to neutralize acid.
● Natural buffers like limestone (CaCO₃) help neutralize acidity slowly.

1. Complete and balance the following equations:
(a) HNO₃(aq) + Ca(OH)₂(aq) →
(b) HNO₃(aq) + K₂CO₃(s) →
(d) HC₂H₃O₂(aq) + Al(OH)₃(aq) →
(e) H₃PO₄(aq) + NaHCO₃(aq) →
(f) H₂SO₄(aq) + Ca(HCO₃)₂(aq) →
(g) HClO₃(aq) + CaCO₃(s) →
2. Barium sulfate can be made using two different neutralization reactions, using the
following substances: Ba(OH)₂, Ba(HCO₃)₂, H₂SO₄
Write chemical equations for these two reactions. Each reaction should have only two reactants.

An Introduction to Molecular Elements and Compounds.pdf

More Related Content

Similar to An Introduction to Molecular Elements and Compounds.pdf (20)

More from jhoyvanwilliams2 (15)

Recently uploaded (20)

An Introduction to Molecular Elements and Compounds.pdf