11-01: Matter, Trends, & Bonding

Ionic Compounds

• Ionic solids are hard, brittle crystals

• Their melting & boiling points are very high

• They conduct electricity only when melted or dissolved

• Otherwise ions are stuck in place

• Ionic bonds form when the difference in electronegativity (or how strong their pull is) between atoms is greater than 1.7

• Generally this means that metals & nonmetals will bond this way with each other

• You can subtract electronegativities and if it's greater than 1.7 (the difference), then it’s ionic

• Properties of ionic compounds are the result of how the ions bond

• Ions form by giving & receiving electrons & produce a rigid lattice (repeating pattern) of strong bonds

• This lattice structure does not allow electrons or ions to move, which therefore creates these brittle, hard solids that do not conduct electricity

• Once an ionic compound melts or dissolves, the ions are then free to move & conduct electricity

• They become free from each other and still have their charge, hence they still conduct electricity

• Binary ionic compounds are composed of only 2 elements (eg. NaCl, Sodium chloride)

• Polyatomic ionic compounds have more than 2 elements (eg. Na2SO4, sodium sulfate)

• Some ions have more than one possible charge (multivalent)

• The IUPAC system shows the charge as roman numbers in the name, and the traditional method uses the suffix “ic” for the higher charge and “ous” for the lower charge

• Mnemonic: i is high, o is low

• Binary compounds are written with the cation first and then named by changing the ending of the second element to “___ide” (always put the metal first)

Eg.

IUPAC: Lead IV bromide

Traditional: Plumbic bromide

Formula: PbBr4 (crossed over, Pb has a charge of 4 and Br has a charge of 1)

• Hydrated salts are ionic compounds that have water molecules attached to them
• The number of water molecules is given with a number in the formula and a prefix in the name
• Eg. magnesium sulfate heptahydrate = MgSO4 • 7H2O

Prefixes:

• Mono = 1
• Di = 2
• Tri = 3
• Tetra = 4
• Penta = 5
• Hexa = 6
• Hepta = 7
• Octa = 8
• Nona = 9
• Deca = 10

Polyatomic Ions

• Some covalently bonded groups of atoms can have a positive or negative charge
• These polyatomic ions can bond and form compounds with properties, just like any other ion can

Common polyatomic ions:

• Nitrate → NO3⁻

• Sulfate → SO4²⁻

• Phosphate → PO4³⁻

• Carbonate → CO3²⁻

• Chlorate → ClO3⁻

• Hydroxide → OH⁻

• Ammonium → NH4⁺

• Elements in the same group of the periodic table have “parent” polyatomic ions with the same number of oxygen atoms and charge (follows the same pattern as “parent”)

Bonded with other elements:

e.g.:

• Magnesium nitrate → Mg(NO3)2

• Calcium sulfate → CaSO4

• Iron II phosphate → Fe(PO4)3

• Manganese IV carbonate → Mn(CO3)2

• Cupric chlorate (Copper II chlorate)  → Cu(ClO3)2

• Lead II hydroxide → Pb(OH)2

• Ammonium sulfide → (NH4)2S

  • Still following the cross over method with the charges
  • Some polyatomic ions have oxygen in them and the number of oxygens can vary

• Polyatomic ions with oxygen can exist with variations in the number of oxygens

• ^^The charge stays the same^^

Rules for polyatomic ions

• Parent (x) → _ate (“default”)

• Add an oxygen (x+1) → per_ate

• Remove an oxygen (x-1)→ ____ite

• Remove 2 oxygens (x-2) → hypoite

• Hydrogens attached to polyatomic ions create compounds that are oxyacids when dissolved in water

• All acids must be labeled (aq) to show that they are dissolved in water, or are aqueous

• Not everything labelled (aq) is an acid, it simply means dissolved in water. When a Hydrogen is present then it is an acid!!

• ==Oxyacids can be named by changing the “ate” polyatomic ions to “ic acid” and “ite” ions to “__ous acid”==

Eg.

NO3⁻ –––hydrogens attach to polyatomic ions making oxyacids→ HNO3 (aq) = nitric acid

NO2⁻ –––hydrogens attach to polyatomic ions making oxyacids→ HNO2 (aq) = nitrous acid

• ==Oxyanions with a charge of -2 or -3 can have hydrogens added and still be charged==
• These ions can bond as any other polyatomic ion can, and include hydrogen in the name

Eg.

CO3²⁻ → HCO3⁻

Carbonate → hydrocarbonate

• Binary acids have 2 elements, and one of them is hydrogen
• Always named hydro_ic acid
• Eg. HCl(aq) is hydrochloric acid
• Organic acids are molecules with COOH as part of their structure
• Acetate ion CH3COO⁻     CH3COOH(aq) is acetic acid (aka vinegar) once it gets the extra Hydrogen bonded to it and becomes aqueous

Atomic Structure

• Atoms are made of 3 basic particles - protons, neutrons, and electrons
• These particles are what makes up the difference between different atoms, ions, and isotopes
• Bohr-Rutherford diagrams and/or Lewis dot symbols are used to show the number of subatomic particles in atoms and help to understand bonding

Ions:

• Ions form when electrons are added/removed to produce a complete outer shell
• Multiple charges are possible with atoms that possess a more complex (complicated) electron configuration (the way that the electrons are arranged)

Isotopes:

• Isotopes are atoms of the same element with different numbers of neutrons

• They can be separated based on mass using a mass spectrometer (the sample must be vapourized, and the electrons are stripped so they get charged, which are brought to emit a beam of particles which then pass through a magnet - heavier particles are deflected less and lighter particles are deflected more. This can determine how heavy/light a particle is but also can determine the number of particles, which is where we get these numbers from)

• The average atomic mass of an element can be determined by using a mass spectrometer.

• The data that is needed to calculate the average atomic mass is the mass and the natural abundance of each isotope

• Since these abundances are constant in nature, a weighted average is calculated. This average is what is recorded on the periodic table

  Find Weighted Average:

• The nucleus of some isotopes is unstable and will decay

• These are called radioisotopes

• Different types of radiation can be released including:

Alpha Particles

•  - Alpha particles: 2 protons and 2 neutrons in the particle are ejected. Always a Helium nuclei

Beta Particles

• - Beta particles: one electron being emitted from the nucleus

Gamma Particles

• - Gamma: photon of electromagnetic radiation (often accompanies alpha or beta → can be added in)

• Radioisotopes have an array of uses like medicine and smoke detectors

• Smoke detectors’ function: smoke gets in and interferes with alpha particles, thus interfering with the circuit and so the alarm goes off

• It is a complete circuit, part of which is completed by alpha particles, smoke breaks the circuit and trips the noisemaker off

Example of Alpha Decay

The alpha decay of Uranium-238

Example of Beta Decay

The beta decay of Carbon-14

Periodic Table Trends

• Groups are vertical columns

• Periods are horizontal rows

• Groups (columns) of elements have similar properties

• Coulomb's law describes the force of attraction or repulsion between 2 charges or objects, with regard to the magnitude of the charge and the distance between the particles

• This law helps us to understand the patterns of properties seen on the periodic table

Force of attraction or repulsion:

• This makes sense of how protons and electrons stay together

• Charge 1 would be the nucleus and charge 2 would be the electrons - distance is how far the electrons are from the nucleus

• Atomic radius is the distance between the nucleus and the valence shell of electrons (size of the atom)

• Atomic radius increases down a group as shells are added

• Atomic radius decreases across a period as more protons are added in the nucleus

• Ionic radius is the same as atomic radius, but for ions

• When an anion (negative) forms, the radius increases because extra electrons repel each other

• When cations (positive) form, the radius decreases because an electron shell is lost

Increasing radius: P, Mg, Cl, Al

Increasing ionic radius: Ca2+, S2-, Na+, F-

If they have the same number of electrons, but protons are different, then the ones with more protons are smaller because then it will have a greater attraction to the nucleus

• First ionization energy is the the amount of energy needed to remove the first electron from an atom (harder to remove from a full valence shell)

• The closer that electron is to the nucleus, the more energy it will take to remove it

• Subsequent ionization energies will require more energy to remove more electrons with a large increase when removing an electron from a new layer

• Electron affinity is the energy released when electrons are added to an atom

• This trend follows the same pattern as first ionization energy

• Electronegativity is a measure of the attraction of atoms for shared electrons

• It’s measured in units called Paulings

• This measure also has the same trend as first ionization energy

• Finding difference in electronegativity ⇒ subtract atoms’ electronegativities

• Reactivity also follows general trends

• Nonmetals tend to react by gaining electrons so they become more reactive up a group, and also more reactive across a period

• Metals react by giving up electrons and are thus more reactive down a group because electrons are less tightly held. Metals are less reactive across a period because the electrons are also more tightly held

Molecules

• Molecules are held together by sharing electrons between atoms with similar electronegativity

• These are mostly nonmetals

• Pure covalent bonds share their electrons equally, and make nonpolar molecules

• Covalent bonds can be polar or nonpolar, depending on the difference between the electronegativities of their atoms

• Electronegativity difference of 0-0.4 ⇒ pure covalent

• Electronegativity difference of 0.4-1.7 ⇒ polar covalent

• Electronegativity difference of 1.7 or greater ⇒ ionic

  

Eg.

HCl has polar bonds ⇒ ∆EN = 3.0 - 2.1 = 0.9

CH₄ has nonpolar bonds ⇒ ∆EN = 2.5 - 2.1 = 0.4

Polarity of molecules depends on the shape of the molecule as well as the polarity of their bonds

• Most covalent bonds are the result of sharing unpaired electrons in the valence shell
• Lewis structures are useful to represent covalent bonding of molecules
• Bond atoms until the valence shell is completed
• This is usually 8 valence electrons, but there are some different atoms that are stable without the full valence shell

Prefix system

• A prefix system is used for naming binary molecules

• Each element gets the prefix for the quantity present, and “___ide” is placed at the end

• The “mono__” prefix gets left off for the first element

• Mono - one

• Di - two

• Tri - three

• Tetra - four

• Penta - five

• Hexa - six

• Hepta - seven

• Octa - eight

• Nona - nine

• Deca - ten

• Bonds between molecules are weaker than bonds between ions in a solid

• This gives molecules distinct properties, like:

• Molecules can be solid, liquid, or gas under standard conditions. Ionic compounds are all solids

• Molecules can be hard or soft. Ionic compounds are all hard

• Molecular compounds’ melting and boiling points are lower than ionic compounds’

• Conductivity of molecules is poor since they do not have a charge or free electrons

Common molecular compounds to note:

• Water → H₂O
• Methane → CH₄
• Ammonia → NH₃
• Vinegar → CH₃COOH
• Carbon dioxide → CO₂
• Glucose → C₆H₁₂O₆

Intermolecular Forces

• Covalent and ionic bonds are the strong attractions that hold atoms together within compounds

• Intermolecular forces are weaker and only occur between molecules

• Joannes Van Der Waal studied factors that affect condensation of molecules

• The forces he studied all hold molecules together with a range of strength but are all weaker than covalent or ionic bonds

• Nonpolar molecules are those where the atoms have similar electronegativity and/or a symmetrical shape

• Polar molecules have permanently charged ends due to differences in electronegativity and asymmetrical shape

Dispersion

• Dispersion forces are the weakest force of attraction that holds molecules together
• This force is the result of the movement of electrons that cause temporary dipoles (charged sides) on the molecule
• These are also called “London Dispersion Forces”
• Dispersion forces are what holds nonpolar molecules together
• The force is very weak in small molecules but increases with molecule size due to there being more electrons to produce these temporary dipoles

Dipole Dipole Attraction

• Polar molecules have permanent dipoles and the oppositely charged parts are attracted to each other
• This attractive force is stronger than dispersion forces but does not increase with molecule size

Hydrogen Bonding

• Hydrogen bonding is the strongest form of dipole dipole attraction

• It occurs when hydrogen is bonded to oxygen (O), nitrogen (N), or fluorine (F)

• The hydrogen is strongly attracted to the lone pairs on adjacent molecules

• Water is a special molecule due to its polarity and very strong hydrogen bonding

• Nonpolar molecules do not dissolve in water while polar ones do

• Polar molecules will dissolve in water as will many ionic compounds

• Water is sometimes called the universal solvent

Crystal Solids

• Solids with a uniform pattern of particles and bonds throughout are called crystalline solids
• Amorphous solids are bonded as well but more randomly

Nonpolar Molecules

• Nonpolar molecules are those with atoms of similar electronegativity and/or a symmetrical shape

This includes:

• Molecules with Carbon and Hydrogen are nonpolar ⇒ eg. C₂H₆

• Molecules with the same element (like diatomic elements) ⇒ eg. Cl₂, Br₂

• Symmetrical molecules ⇒ eg. CH₄, CCl₄, CO₂

• Nonpolar molecules form crystals by dispersion force attraction

• The melting point and boiling points are low but increase with molecular size

• These compounds have low conductivity

• Not water soluble, but can be dissolved in nonpolar solvents

Polar Molecules

• Polar molecules have atoms with different electronegativity and asymmetrical shape

This includes:

• Molecules with Hydrogen bonded to O, N, F

• Molecules with 2 atoms of very different electronegativities

• Asymmetrical molecules (eg. H₂O, C₂H₅OH)

• Polar molecules bond together by dipole-dipole or hydrogen bonding

• These crystals will have higher melting and boiling points than nonpolar crystals

• These compounds have low conductivity

• They are soluble in water, but not in nonpolar solvents

Ionic Solids

• Ionic solids form from a repeating pattern of oppositely charged ions

• These solids have high melting points due to strong bonds

• The greater the charge difference, the stronger the bond

• They are also brittle and hard

• They don't conduct electricity because their ions cannot move

• Ionic crystals will only conduct electricity when they're either melted or in a solution

• The ions must be free in order to conduct a charge

Covalent Network Crystals

• Covalent network crystals are held together with covalent bonds throughout their structure

• This makes them the hardest substance known

• They have very high melting points

• Different forms of covalent network are called allotropes

  

^^Different ways for carbon to be bonded → crystalline = where all bonds are the same and amorphous = there is a range, they're not all the same

Metallic Crystals

• Metallic crystals form by donating electrons to a “sea” shared between cations throughout the entire solid

• Metals melt over a range of temperatures

• Free electrons make metals malleable, ductile, shiny, and very conductive of electricity and heat

  

• Electrons hold cations together but not in a specific way

• This is why metals have all the traits they do (for the most part), because the electrons bounce around