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

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• 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

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• 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

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• 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

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• 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)

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• 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

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• 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)

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• 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

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Prefixes:

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

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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

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Common polyatomic ions:

• Nitrate → NO3⁻
• Sulfate → SO4²⁻
• Phosphate → PO4³⁻
• Carbonate → CO3²⁻
• Chlorate → ClO3⁻
• Hydroxide → OH⁻
• Ammonium → NH4⁺

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• 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”)

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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

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• Polyatomic ions with oxygen can exist with variations in the number of oxygens
• ^^The charge stays the same^^

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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) → hypo______ite

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• 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

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• ==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

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• 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

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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:

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• The nucleus of some isotopes is unstable and will decay
• These are called radioisotopes
• Different types of radiation can be released including:

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Alpha Particles

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•  - 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)

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• 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

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Example of Alpha Decay

The alpha decay of Uranium-238

Example of Beta Decay

The beta decay of Carbon-14

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Periodic Table Trends

• Groups are vertical columns
• Periods are horizontal rows
• Groups (columns) of elements have similar properties

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• 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

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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

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• 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

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• 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

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• 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

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• Electron affinity is the energy released when electrons are added to an atom
• This trend follows the same pattern as first ionization energy

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• 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

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• 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

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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

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• 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

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• 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

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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

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• 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

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Common molecular compounds to note:

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

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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

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• 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

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• 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

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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

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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

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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

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• 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

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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

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• 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

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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)

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• 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

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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

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• They are also brittle and hard
• They don't conduct electricity because their ions cannot move

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• 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

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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

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\ ^^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

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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

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• 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

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