Chemistry Unit 4

Chemical Bond

the attraction between atoms as a result of transfer or sharing of valence electrons.

Octet Rule

Chemical compounds tend to forms so that each atom, by gaining, losing, or sharing electrons, has an “octet” of electrons in its highest occupied energy level (having a full s and p sublevel). There are exceptions – hydrogen. When the highest occupied energy level (valence shell) is filled with electrons, the atom is stable and not likely to react like their noble gases

Ionic Bonds

Valence electrons are transferred from one atom to another. Ionic bonds are the force that holds cations and anions together – formed between a METAL and a NONMETAL. Properties include: High melting points, conduct electricity in the molten state (electrolyte), soluble in water, and tend to be hard and brittle due to strong bonding

Ionization Energy

the amount of energy used to remove an electron from an atom.

-left to right: strength of energy generally increases -bottom to top: strength of energy generally increases Francium vs. Fluorine

Chemical Formula

Notation that shows what elements a compound contains and the ratio of atoms or ions of these elements in the compound

ex.) NaCl – sodium chloride = 1 sodium and 1 chloride

Oxidation Numbers

the number associated with how many electrons an atom wishes to lose or gain in an ion bond; +1,+2,+3,+4&-4, -3,-2,-1

Covalent Bonding

Valence electrons are SHARED to complete the valence shell atom. These bonds occur between 2 nonmetals. Note:Carbon and Hydrogen always bond covalently! Properties:Low melting points (ex. sugar), doesn’t conduct electricity, brittle


Neutral group of atoms joined together by one or more covalent bonds. The attraction between the shared electrons and protons in each nucleus hold the atoms together in a molecule.

Multiple Bonds

Double Bond:when 2 atoms share 2 pairs of electrons Triple Bond:when 3 atoms share 3 pairs of electrons

Metallic Bonds

Attraction between a metal cation and the shared electrons that surround it. In metals, valence electrons are free to move among the atoms creating a pool of “shared electrons”

Attracting Forces

The cations in a metal form a lattice that is held in place by strong metallic bonds between the cations and the surrounding valence electrons. Electrons are moving among atoms. The number of electrons don’t change so the metal is neutral.

Bond Strength

The more valence electrons in the “pool”, the stronger the metallic bonds will be. Bonds in an alkali metal are weak because they contribute only one electron. Transition metals have more valence electrons to contribute, therefore they have stronger bonds and are harder


the relative tendency of an atom to attract electrons to itself when it is bonded to another atom

Electronegativity Difference (E.D.)

– e.d. greater than 1.7 is ionic(greater e.d. = more ionic properties) – e.d. less than 1.7 is covalent To determine e.d., subtract the smaller electronegativity value from the larger electronegativity value

*There are always exceptions! (The best way to identify ionic or covalent is by looking at what is bonded together)

Chemical Formula

chemical symbols that represent the composition of a compound


number which represents “ratio” of elements in a compound

Empirical Formula

“simplest” whole number ratio in which atoms combine to form a compound ex.)C6H12 = molecular formula; CH2 = empirical formula

*Ionic compounds are always represented by empirical formulas

Molecular Formula

the total number of atoms of each element in one molecule

Rules for Writing Ionic Formula

Write down symbol and charge of cation. Write down symbol and charge of anion. Drop the “charges (+ and -)” and “criss-cross”(the charge number). Use parenthesis around polyatomic ions when new subscript is used (other than 1)

Variable Charge Ions (the transition metals)

“Stock System” uses Roman Numerals to identify charge of those cations with variable charges.

*Zinc(Zn) ALWAYS has a +2 charge while Silver(Ag) ALWAYS has a +1 charge

Naming Ionic Compounds

Cation(+) name always goes first. Anion(-) name always goes second. Use “stock system”(roman numerals) to identify charge of variable charge cations(all transitional metals except Zn & Ag). Monatomic anions(found on the periodic table)end with “-ide”.

ex.)MgO: magnesium + oxygen = magnesium oxide Polyatomic ions retain their endings.

ex.)Na3PO4: sodium + phosphate = sodium phosphate

Naming Covalent Compounds

Covalent compounds are formed between 2 NONMETALS. First element is named using full name. Second element is named as if it were an anion(ends in -ide). Prefixes are used to denote the number of atoms present. The prefix “mono” is NEVER used for naming the first element, but it can be used for naming the second element.

ex.)NO = nitrogen monoxide; P2O5 = diphosphorus pentoxide

# of Atoms prefixes











Rules for Writing Covalent Formula

If no prefix is on the first element name, assume there is only one. If there is a prefix on the first element name, write the symbol and the subscript for that prefix. There will ALWAYS be a prefix on the second element name! Ex.)Carbon dioxide = CO2; Diphosphorus tetracholoride = P2Cl4

Lewis Structures

1. Ionic Compounds

Uses electron dot diagrams to represent compounds. Shows how atoms are gaining or losing electrons to form ionic bonds.

1. For ionic compounds, represent cations and anions with their charges. Cations – write the symbol and charge; NO VALENCE ELECTRONS ARE DRAWN. Anions – place the dot diagram, with 8 valence electrons, in brackets with the charge on the outside. If there is more than 1 ion in the compound, put the coefficient in front of the bracketed anion

Ex.)Na2S: 2Na+1 [ [image] ]-2
Magnesium chloride: MgCl2 : Mg+22[image]-1

2. For covalent compounds – a bond will form between electrons that are to be shared between atoms. An atom will bond in places with only 1 valence electron, or one “dot”. When drawing dot diagrams for covalent compounds, you connect single bonds with a line, double bonds with 2 lines and triple bonds with 3 lines. Place all other non-bonding electrons around their respective atoms.

Ex.) Water: [image]

Nitrogen triflouride:[image]

C2H5Cl: [image]


uneven charge distribution (due to electrons being shifted to one side). The element with the highest electronegativity will “pull” the electrons closer

Electronegativity Difference w/Polarity

if e.d. is greater than or equal to 1.7 = ionic if e.d. is less than 1.7 = covalent e.d.=0.0 to 0.3: nonpolar e.d.=0.31 and up: polar

*the degree of polarity becomes greater as electronegativity difference increases

Polar covalent

covalent bond in which the electrons are not shared equally resulting in an unbalanced distribution of charge. -When atoms form a polar covalent bond, the atom with the greater attraction for electrons has a partial negative charge charge. The other atom has a partial positive charge

Nonpolar covalent

covalent bond in which the bonding electrons are shared equally resulting in a balanced distribution of charge

Polarity in Molecules

to determine if a molecule is polar of non-polar, you must determine: -the polarity of the individual bonds in the molecule. If all bonds are non-polar, then the whole molecule is nonpolar regardless of its shape (a rare case). -the shape or geometry of the molecule. If there is symmetry in the molecule, than it is nonpolar & vice versa. Once you know which bonds in the molecule are polar and non-polar, you must use the shape of the molecule – if 2 polar bonds are oriented correctly, they can cancel each other out (like 2 equally strong people pulling in opposite directions on a rope – the robe won’t move)

1.Intramolecular Bonds

2.Intermolecular Bonds

1.the bonds holding the molecule itself together. The prefix intra- comes from the Latin word meaning “within or inside” ex.)The covalent bonds between hydrogen and oxygen atoms in a water molecule

2.The bonds between molecules that hold them all together. The prefix inter- comes from the Latin word meaning “between”

Bond Strength

Ionic bonds are stronger than covalent bonds. Ionic bonds typically occur in large quantities whereas covalent bonds typically form in isolated molecules. Ionic bonds are attractions between a positive and negative(there is a physical transfer of electrons rather than just sharing), which tends to be stronger. Polar bonds are stronger than non-polar – because they have an unequal distribution of electrons, one atom has a higher attraction for the electrons. Intermolecular forces have different strengths, but are overall weaker than intermolecular forces.

-Boiling points: the higher the boiling point, the stronger the forces between particles

-Bond Strength in general order: ionic, covalent, intramolecular, intermolecular


1.Dipole-dipole forces

equal but opposite charges separated by a short distance. From positive to negative. Arrows point toward the negative and crossed end sits at the positive.

1.The attraction between polar molecules (partial positive of one molecule is attracted to the partial negative of another molecule. This force acts only on nearby molecules. (The strongest intermolecular forces exist between polar molecules.

-Act as dipoles(unequal charge distribution)) Polarities in dipoles can be increased or decreased depending on the orientation and direction of dipoles


Hydrogen Bonding

Occurs when a hydrogen that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule. Occurs in compounds containing H-F, H-O, H-N bonds. (Highly electronegative differences between hydrogen and fluorine, hydrogen and oxygen, and hydrogen and nitrogen create highly polar bonds. Hydrogen is so small, allowing the atom to get close to an unshared pair of electrons on another molecule

London Dispersion Forces

Intermolecular attractions from the constant motion of electrons and instantaneous and temporary dipoles. In atoms and molecules, electrons are always in motion. At any instant, the electron distribution within the atom or molecule may be uneven(more electrons on one side vs. the other for maybe a nanosecond). Even though it may be momentary, the uneven distribution creates a positive charge on one side and a negative charge on the other, causing a temporary dipole. If another neighboring atom or molecules has the same temporary dipole, they may be attracted to each other. London dispersion forces act between all atoms and molecules, but is the only intermolecular force in noble gas and non polar molecules

Atomic Radius


2.How does atomic radii affect electronegativity?

The radius of an atom without regard to surrounding atoms.

1.Atomic radii increases from top to bottom and right to left on the periodic table. Trends occur because: -increase in energy level -increase in # of protons and # of electrons -decrease in electrostatic attraction; increase in atomic size

2.Electronegativity increases as size of atomic radius decreases (The protons in a nucleus of an atom that has a large atomic radius will have more difficulty than an atom with a small radius because the farther away the nucleus is from the electrons, the less attraction(electronegativity) it will have

Ionic Size – would the radius of an ion be larger or smaller than the atom it comes from?

Cations lose electrons and become smaller – they drop down in number of energy levels Anions gain electrons and become larger – repulsion between electrons increases so electrons expand the cloud to accommodate

Molecular Shape


this is due to the fact that the electron “pairs” repel each other as much as possible whether they are in a bond or not.

1.V:alence S:hell E:lectron P:air R:epusion Shared pair of electrons – involved in bond Unshared pair of electrons – not involved in bond

Diatomic Molecules

Molecules composed of only 2 types of atoms

Ex.)H2, HCl These molecules must be linear because they consist of only 2 atoms To predict shapes of more complicated molecules, you have to consider the location of all electron pairs surrounding the bonded atoms (VSEPR method)

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