Unit 3- Bonding

Metallic Character of Elements

Metallic Character of Elements

• The weaker an atom attracts electrons (Z_eff) the more metallic it is

Three categories for Metallic Character of the Elements

metals, nonmetals and metalloids

Metalloids

have properties somewhere between metals and nonmetals, though they are a generally bit more like nonmetals.

Covalent bonds

a chemical bond that involves the sharing of electrons to form electron pairs between atoms

Ionic bonds

a type of chemical bond that occurs between two atoms when one atom transfers electrons to another atom. It is the electrostatic attractions between positive and negative ions

• Ionic compounds are traditionally thought of as forming from the reaction of metals and nonmetals (though they form other ways as well)

cations

positive ions (removing electrons)

anions

negative ions (gaining electrons)

electrostatic attraction

attraction between positive and negative charges

Noble Gases and Bonding

only elements that exist as single, un-bonded atoms. This is due to their valence shell being full & stable.

Covalent Bonding Principles

electrons are shared between atoms to achieve a stable configuration

• Covalent bonding usually results in atoms obtaining an octet of valence electrons through a combination of shared and unshared valence electrons (WITH EXCEPTIONS OF HYDROGEN, which only has 1 Principle energy level therefore only needing to complete shell with 2 electrons)

Bonding Pairs

shared electron pairs between atoms. Belongs to two or more atoms.

Non-bonding pairs/lone pairs

unshared pairs of valence electrons. Only belong to one atom.

ionic compounds

empirical formula where ratio is always the lowest whole number mole ratio of elements in the compound.

oxidation

an increase in oxidation state (charge) caused by the loss of electron(s)/electron ownership

reduction

a decrease in oxidation state (charge) caused by the gaining of electron(s)/electron ownership

covalent triple & double bonds

only elements that exist as single, un-bonded atoms. This is due to their valence shell being full & stable.

Ionic Chemical Formulas & Names

Determining the formula of an ionic compound is done by balancing the ion charges so that the overall charge is neutral.

• If the positive and negative charges on the ions are not equal, then subscripts are added to balance the charges. The “Criss Cross Apple Sauce” Method is a simple way to do this.

Bonding Pairs

shared electron pairs between atoms. Belongs to two or more atoms.

Polyatomic Ions

ions made up of more than one atom. These atoms are generally nonmetals and are covalently bonded together, but have an overall positive or negative charge and therefore act like ions.

• They have their own specific formulas, charges and names that must be memorized.

Non-bonding pairs/lone pairs

unshared pairs of valence electrons. Only belong to one atom.

Lewis Dot Structures

simplified models for representing the covalent bonding between atoms, where each electron is represented by dots. Electrons are shown in pairs, with bonding pairs usually being represented as a line segment.

Diatomic elements

Diatomic elements are molecules composed of two atoms of the same element bonded together. Examples of diatomic elements include oxygen (O2), nitrogen (N2), and hydrogen (H2).

Lewis dot structures for phosphorus (P4) and Sulfur (S8)

Ionic Chemical Formulas & Names

Determining the formula of an ionic compound is done by balancing the ion charges so that the overall charge is neutral.

• If the positive and negative charges on the ions are not equal, then subscripts are added to balance the charges. The “Criss Cross Apple Sauce” Method is a simple way to do this.

Rules for Naming Simple Ionic Compounds

Name both ions, cation (atom with positive charge) first.

• Normal metal cations are named like the element: calcium → calcium

• Normal non-metal anions are named with an “-ide” suffix:

  chlorine → chloride

  oxygen → oxide

  hydrogen → hydride

  nitrogen → nitride

Naming Ionic Compounds

1. Name the metal (the cation) as it appears on the Periodic Table.

2. For the non-metal (the anion) write the name on the Periodic Table and then replace the ending with ide.

CaCI2 = Calcium chlorine = Calcium chloride

AlN= Aluminum nitrogen = Aluminum nitride

3. Use the total charge on the non-metal (or polyatomic ion) find the charge on the transition metal.

Diatomic elements

Diatomic elements are molecules composed of two atoms of the same element bonded together. Examples of diatomic elements include oxygen (O2), nitrogen (N2), and hydrogen (H2).

delocalised electrons & metallic bonding

• In a molecule, ion, or solid metal, electrons that are not associated with a single atom or covalent bond are referred to as "free electrons". In the elemental state, when there are no other elements present to accept the electrons and form an ionic compound, the outer electrons are loosely held by the metal atom's nucleus and have a tendency to move around. These free electrons are not fixed in one position and can spread themselves throughout the metal structure. This allows them to move freely within a regular lattice structure, forming cations in the process.

  (think about it like a close neighbourhood of families where the children do not belong specifically to any one set of parents but are free to wonder between the homes, causing a close association between the families)

Characteristic physical properties of metals

• good electrical conductivity (electrical circuits use copper)

• good thermal conductivity (pots and pans used for cooking)

• malleable, can be shaped under pressure (moulded into many forms including machinery and structural components of building and vehicles)

• ductile, can be drawn out into threads (electric wires and cables)

• high melting points (high speed tools and turbine engines)

• shiny, lustrous appearance (ornamental structures and jewellery)

alloy

An alloy is a mixture of two or more metals or a metal and a non-metal. It has different properties than the individual elements. Alloys are commonly used in manufacturing to create materials with specific properties, such as increased strength or resistance to corrosion. Examples include brass, steel, and bronze.

strength of the metallic bond is determined by

• the number of delocalided electrons

• the charge on the cation

• the radius of the cation

melting points of metals

the more delocalised electrons present and the smaller the radius of the atom, the higher the melting point of the metal.

strength of metallic bonding

decreases down a group as the size of the cation increasing, reducing the attraction between the delocalised electrons and the positively charged protons in the nucleus

metallic bonding of metals (trends)

Left to right across a period

• Increasing melting point

• greater attraction between ions and delocalised electrons

• lower degree of reactivity

Down a group

• decreasing melting point

• weaker attraction between ions and delocalised electrons

• higher degree of reactivity

Alloy structure

as an alloy mixture solidifies, ions of the different metals are scattered through the lattice, forming a structure of uniform composition.

• Alloys contain metallic bonds as the delocalised electrons bind the lattice. The production of alloys is possible because of the non-directional nature of the delocalised electrons and the fact that the lattice can accommodate ions of different size.

Alloy properties

they have properties that are distinct from their component elements due to the different packing of the cations in the lattice.

• The regular arrangement of atoms in a pure metal is interrupted in the alloy by the presence of different cations , making it more difficult for atoms to slip over each other and so change the shape.

• The alloy is often stronger, more chemically stable, and more resistant t corrosion than its component elements. For example, steel, which is an alloy of iron is 1000 times stronger than iron.

Common alloys

Alkanes

simplest class (family) of organic compounds. They are hydrocarbons in which the carbon atoms are held together by single bonds.

General formula: C_nH_2n+2

IUPAC Naming

Root/Parent Name: the number of carbon atoms in the longest continuous chain

• Meth

• Eth

• Prop

• But

• Pent

• Hex

Suffix: The class of compound

• ane: alkane (no functional group)

Alkenes

• General formula: C_nH_2n

• Functional Group: C=C

• Naming Suffix: -ene

Alkynes

• General formula: C_nH_2n-2

• Functional Group: N/A

• Naming Suffix: -yne

Alcohols

• General formula: C_nH_2n+1OH

• Functional Group: OH

• Naming Suffix: -anol

Lewis dot structures for phosphorus (P4) and Sulfur (S8)

Noble Gases and Bonding

• double covalent bond is covalent bond formed by atoms that share two pairs of electrons.

• triple covalent bond is a covalent bond formed by atoms that share three pairs of electrons

Bonding Pairs

shared electron pairs between atoms. Belongs to two or more atoms.

Non-bonding pairs/lone pairs

unshared pairs of valence electrons. Only belong to one atom.

Transition Metal

element with an incomplete d-sublevel (1-9 d electrons) in one or more of its oxidation states (Oxidation states can be more or less be thought of as charges the transition can have as ion or within compounds)

What metal in the d-block is not a transition metal?

Zinc, as it has a complete d-block configuration for all of its oxidation states

What is the only oxidation state of copper which allows it to be a transition metal?

When it forms a Cu²⁺ oxidation state- electron configuration is [Ar]3d⁹, meaning that it has an incomplete d-block configuration

Metallic bond strength & physical properties

• Both d & s electrons can become delocalised and contribute to the metallic bonding in transition metals. This makes the metallic bonds in transition metals much stronger than those in typical s & p block metals. As a result, transition metals typically have much different properties than s & p block metals, listed below:

  High electrical and thermal conductivity

  High melting points

  Malleable & Ductile

  High tensile strength

Why does Zeff remain (almost) constant along transition elements?

Occurs due to the similar shielding effect of inner electrons. The inner electrons shield the outer electrons from the full positive charge of the nucleus. As we move across a period in the periodic table, the number of protons in the nucleus increases, which would suggest an increase in Zeff. However, the number of inner electrons also increases, providing more shielding and offsetting the increase in Zeff.

Transition Metals Important properties

• Multiple oxidation states

• Catalytic behaviour

• Show Magnetism

• Form a variety of complex ions

• Form coloured compounds

• High electrical conductivity (due to more delocalised electrons, causing increase of strength of metallic bonds)

• High melting points (due to more delocalised electrons, causing increase of strength of metallic bonds)

Transition Metals and multiple oxidation states

• Metals in group 1, 2, 3 form ions with only one charge

• Every transition metal forms a +2 ion, which is due to it losing its 2 valence electrons (as all transition metals have s-orbitals as their valence electrons)

• As the s-electrons and inner d-electrons are high in energy (due to convergence), it is possible for not just the outer s-electrons to be removed, but also the d-electrons. This explains how multiple oxidation states can be formed.

Naming with Transition metals

• oxidation state must be included in the name as a roman numeral (oxidation number) (e.g Copper (II) oxide, CuO)

• nonmetal in the compound will have its ending switched to ‘ide’, while metal will maintain its regular name (e.g Titanium (II) Oxide)

Transition Metals: Catalytic Behaviour

can act as Catalysts

• Transition metals are particularly effective catalysts because they have the ability to change their oxidation state during a reaction, which allows them to participate in multiple steps of the reaction.

Catalysts

a substance that speeds up a chemical reaction, or lowers the temperature or pressure needed to start one, without itself being consumed during the reaction.

Transition Metals: Magnetism

• Spinning electric charges create magnetic fields. Each electron, therefore, creates its own very tiny magnetic field. With most substances, the electrons are arranged randomly and these magnetic fields cancel out, so there is no net magnetic field.

If subjected to an external magnetic field, the electrons within an element will align with, or against, the applied magnetic field. Depending on the electron configuration of an element, it will respond differently to these applied magnetic fields.

• Magnetism is related to unpaired electrons

Giant Covalent Solids & Structures

while most covalently bonded substances are molecular, there are a few substances that form a different kind of structure called giant covalent structures.

• always solids at room temperature

• strongest bond

• much harder that molecular solids

• higher melting points than molecular solids

Typical example of a giant covalent solid is a diamond, which is made of carbon. Other examples are of silicon (Si) and silicon dioxide (SiO2)

Covalent Molecular Bonds

the sharing of electron pairs between atoms that form molecules, with weak intermolecular forces between the molecules.

• can break apart easily when heated (low melting point)

• forms weakest bond

Physical Properties of all Bonds

What is hardness determined by?

depends on strength of the attraction between neighbouring particles. Stronger inter particle attractions lead to solids that are harder, and in many cases, more brittle, meaning they tend to break-split apart rather than bend.

• Weaker inter-particle attractions can lead to substances being softer solids. If attractions are weak enough than substances will be liquid or gas, and so the term “hardness” wouldn’t apply.

Malleability

the ability to be shaped, which is also linked to hardness of a metal

What is electrical conductivity determined by?

The ability to conduct electricity requires presence of free-moving charges, so that current (a flow of electrons) can occur in a circut.

What is Thermal Conductivity determined by?

based on how efficiently atoms can pass along thermal energy to one another. Atoms that are more tightly packed together tend to be better conductors of thermal energy, while molecular substances, especially gases, tend to do so less efficiently.

What is Melting & Boiling points determined by?

depend on the strength of the attraction between neighbouring particles. Stronger inter-particle attractions lead to higher melting and boiling points

What is Solubility determined by?

determined by how well the particles of one substance will mix with (form attractions with) the particles of another substance. The expression “like dissolves like” is a simple and useful term to remember, and it generally refers to the polarity of substances being mixed.

Lattice Enthalpy

the enthalpy (strength) required to convert 1 mole of an ionic compound into gaseous ions

• can be thought of as a direct way of measuring the strength of ionic bonds

• stronger ionic bonds lead to higher lattice enthalpies (more energy to break the bonds)

Factors affecting Lattice Enthalpies

• More/Higher Charges = Stronger bonds

• Smaller Ionic Radii= Stronger bonds

Bond Enthalpies

The energy required to break 1 mole of a bond in the gas phase.

• Bond breaking involves the separation of particles which are attracted to each other. It is therefore always an endothermic process.

Factors affecting Bond Enthalpies

• More Bonds (Such as Multiple Carbon bonds)= Stronger Bonds

• Smaller Atomic Radii= Stronger Bonds

Ligands

species with at least 1 un-bonded electron pair that form coordinate covalent bonds with a metal ion.

Examples of Ligands: (H₂O), (NH₃), (CO), (CN^-), (Cl^-), (OH^-).

coordinate covalent bonds

a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.

• One electron has DONATED its shared electron pair to the other atom to form a covalent bond.

Formation of Transition Metal Complexes

The positive charge density of the metal ions is so great, it attracts the electron pairs of nearby molecules and ions. This results in the formation of a coordinate covalent bond between the metal ion and the ion/molecule (now called a ligand)

A complex

The large structure formed between the covalently.bonded transition metal ion and ligands. If they have an overall charge, they are called complex ions.

Chemical formulas of complexes & complex ions

Complex ions are usually surrounded by brackets []. The overall charge is written outside the brackets.

e.g:

• [Co(NH₃)₆]³⁺

• [CuCl₄]^2-

Normal complex:

• Fe(H₂O)₃(CN)₃

All Polyatomic Ion Names

+1 CHARGE

• Ammonium

-1 CHARGE

• Hydroxide

• Nitrate

• Hydrogen Carbonate

-2 CHARGE

• Sulfate

• Carbonate

-3 CHARGE

• Phosphate

Ionic Hydrates

• ionic compound that is linked to at least one molecule of H₂O

• Water molecules are attracted to the ions through ion-dipole interactions, which fill gaps in the lattice of ions. This makes the ions HYDRATED, where the H₂O molecules are considered in their structure/chemical formula.

Hydrous

containing water (in Ionic structures)

Eg: CoCl₂•6H₂O

Colbat (II) Chloride 6-hydrate

Anhydrous

not containing water (in Ionic structures)

Eg: CoCl₂

Colbat (II) Chloride

Formula Units

the term used when referring to the number of particles in an ionic compound, since saying “molecules” is inaccurate.

Example:

1 mole of water= 6.022 × 10²³ H₂O molecules

1 mole of sodium chloride= 6.022 × 10²³ NaCl formula units

Determining # of atoms & # of ions in Ionic compound

For atoms:

MgSO₄= 6 atoms

times # of atoms by avogadro’s number

6(6.022 × 10²³)= 3.61 × 10²⁴

For ions:

MgSO₄= 2 ions

times # of atoms by avogadro’s number

2(6.022 × 10²³)= 1.20 × 10²⁴

Empirical formulas of Ionic Hydrates

1. Determine mass of anhydrous (ionic compound without water) ionic compound and the mass of the water

2. Determine # of moles in the anhydrous ionic compound and # of moles of water

3. Divide the number of water moles by the moles of the ionic compound. This gives the number of moles of water per 1 mole of ionic compound, which is the coefficient of water in the chemical formula.