1. Forces that hold atoms together

2.  There are several major types of bonds. Ionic, covalent and metallic bonds are the three most common types of bonds.  Covalent bonds – electrons are shared between atoms.  Ionic bonds – electrons are transferred between atoms, creating cations and anions.  Metallic bonds – two or more metals bonded together.

3.  There are two different types of covalent bonds, polar covalent and nonpolar covalent. ◦ polar covalent – electrons are not shared equally between the two bonded atoms. The electrons are pulled toward the more electronegative of the elements. ◦ nonpolar covalent – electrons are shared equally between the two bonded atoms.

4. Electronegativities

5.  Ionic bonds are formed when there is an electronegativity difference (  EN) greater than 2.0.  Polar covalent bonds form when there is a  EN between 0.5 and 1.7.  Nonpolar covalent bonds form when there is a  EN between 0 and 0.49.

6.  If the  EN is between 1.7 and 2.0, an ionic bond will form if a metal is one of the elements, and a polar covalent bond will form if only nonmetals or metalloids are present.  What type of bond is formed between the following elements?  N and O  K and F  Mg and Cl  P and F  C and H

7.  Metals tend to lose their valence electrons, leaving a complete octet in their next-lowest energy level.  Sodium – (1 valence electron) loses 1 electron and becomes Na +1.  Na ([Ne]3s 1 ) → 1e - + Na +1 ([Ne])  Calcium – (2 valence electrons) loses 2 electrons and becomes Ca +2.  Ca ([Ar]4s 2 ) → 2e - + Ca +2 ([Ar])

8.  Nonmetals tend to gain or share valence electrons to complete an octet in their highest energy level.  Oxygen – (6 valence electrons) gains two electrons to become O -2.  O ([He]2s 2 2p 4 ) + 2e - → O -2 ([He] 2s 2 2p 6 )  Phosphorus – (5 valence electrons) gains three electrons to become P -3.  P ([Ne]3s 2 3p 3 ) + 3e - → P -3 ([Ne] 3s 2 3p 6 )

12.  Ionic bonds – forces of attraction that bind cations and anions together.  Ionic compound – consists of electrically neutral group of ions joined by electrostatic forces.  Example: Sodium chloride

13.  At room temperature, most ionic compounds are crystalline solids, where ions are arranged in various 3-D patterns.  Because of the large attractive forces of the ions to each other the compounds become very stable and have high melting points.

14. Source: ©Clyde H. Smith/Peter Arnold, Inc.

17.  Scientists have learned that all of the elements within each group behave similarly because they have the same number of valence electrons.  Valence electrons - # of electrons in the highest occupied energy level of an atom.  The number of valence electrons is related to the group numbers on the periodic table.

18.  Group 1 elements = 1 valence electron.  Group 2 elements = 2 valence electrons.  Groups 3-12 elements = 2 valence electrons.  Group 13 elements = 3 valence electrons.  Group 14 elements = 4 valence electrons.  Group 15 elements = 5 valence electrons.  Group 16 elements = 6 valence electrons.  Group 17 elements = 7 valence electrons.  Group 18 elements = 8 valence electrons.

19.  1.Multiply the number of valence electrons by the number of moles of each element.  2.Add up all the electrons for each of the elements.  3.If there is a charge and it is negative, add that number of electrons to the total.  4.If there is a charge and it is positive, subtract that number of electrons from the total.  Total # of electrons should always be an even number!

20.  Determine the number of valence electrons in each of the following compounds and ions:  NH 4 +1  CH 2 ClBr  PO 4 -3

21.  Valence electrons are the only electrons involved in bonding, and are the only ones written when drawing electron dot structures.  In forming compounds, atoms tend to achieve the electron configuration of a noble gas, having 8 valence electrons which as known as having a stable octet (octet for 8 valence electrons).

22.  Also known as electron dot symbols  Use symbol of element to represent nucleus and inner electrons  Use dots around the symbol to represent valence electrons  Elements in the same group have the same Lewis symbol ◦ Because they have the same number of valence electrons  Cations have Lewis symbols without valence electrons  Anions have Lewis symbols with 8 valence electrons

23.  Structural formula – chemical formulas that show the arrangement of atoms in molecules and in polyatomic ions.  Octet rule – atoms gain or lose electrons to acquire the stable electron configuration of a noble gas, usually having 8 valence electrons.  Exceptions to the octet rule: ◦ H needs 2 electrons to be stable ◦ Be needs 4 electrons to be stable ◦ B needs 6 electrons to be stable

24.  You can represent the formation of the covalent bond in H 2 as follows: –This uses the Lewis dot symbols for the hydrogen atom and represents the covalent bond by a pair of dots.

25.  The shared electrons in H 2 spend part of the time in the region around each atom. –In this sense, each atom in H 2 has a helium configuration. : HH

26.  The formation of a bond between H and Cl to give an HCl molecule can be represented in a similar way. –Thus, hydrogen has two valence electrons about it (as in He) and Cl has eight valence electrons about it (as in Ar).

27.  Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures. –An electron pair is either a bonding pair (shared between two atoms) or a lone pair (an electron pair that is not shared). bonding pair lone pair

28.  Single covalent bond – a bond in which two atoms share a pair of electrons.  Double covalent bond – a bond in which two atoms share two pairs of electrons.  Triple covalent bond – a bond in which two atoms share three pairs of electrons.

29.  Steps for Drawing Lewis-dot structures 1. Determine the number of valence electrons in the molecule. - When drawing determining valence electrons for an ion, add electrons if it an anion, and subtract electrons if it is a cation. 2. The first element in the compound will be the central atom. Exception: hydrogen will never be the central atom.

30. Steps for Drawing Lewis-dot Structures 3.Use one pair of electrons to bond each outer or terminal atom to the central atom. 4.Make all outer or terminal atoms stable using the valence electrons (8 total dots except for Hydrogen which only needs 2). 5.Put any remaining electrons around the central atom as lone pairs.

31.  Draw the Lewis structure for:  NH 3  PO 4 3-  CHFClBr  PF 5

32.  Single covalent bond – a bond in which two atoms share a pair of electrons.  Double covalent bond – a bond in which two atoms share two pairs of electrons.  Triple covalent bond – a bond in which two atoms share three pairs of electrons.

33.  Single bonds are longer (length between the atoms) than double and triple bonds.  Double bonds are longer than triple bonds.  Single bonds are not as strong as double bonds, and can be broken much easier than double bonds.  Triple bonds are stronger than double bonds.

34.  If you have used up all of the valence electrons and you still need two more electrons to make the central atom stable, you must have one double bond.  If you still need four more electrons to make the central atom stable, you must have either one triple bond or two double bonds.  Double and triple bonds exist most commonly between C, N, O, and S atoms.

35.  Draw Lewis structures for:  NOCl  CO 2 N2N2  SiO 3 -2

36.  The valence-shell electron pair repulsion (VSEPR) model predicts the shapes of molecules and ions by assuming that the valence shell electron pairs are arranged as far from one another as possible. –To predict the relative positions of atoms around a given atom using the VSEPR model, you first note the arrangement of the electron pairs around that central atom.

37.  The following rules and figures will help discern electron pair arrangements. 1.Draw the Lewis structure 2.Determine how many bonding pairs are around the central atom. Count a multiple bond as one pair. 3.Determine how many lone pairs, if any, are around the central atom. All diatomic molecules have a linear shape.

38. 3 pairs Trigonal planar 2 pairs Linear 4 pairs Tetrahedral 5 pairs Trigonal bipyramidal 6 pairs Octahedral

40.  NH 3  PO 4 3-  CHFClBr  CO 2  NOCl N2N2 H2SH2S  SiO 3 -2

41.  Nonpolar covalent bond – equal sharing of electrons between two atoms.  Polar covalent bond – unequal sharing of electrons between two atoms.  In polar covalent bonds the electrons are pulled closer to the atom with the larger electronegativity value.  Polar bonds can create polar or nonpolar molecules and ions.

42.  The easiest way to determine if a molecule or ion is polar or nonpolar is to look at the central atom.  If the central atom has lone pairs of electrons, the molecule or ion is polar.  If the central atom does not have any lone pairs of electrons, the molecule or ion is nonpolar.

43.  SO 2  PO 4 -3 N2N2  BrO 2 -1

44. Attractions Between Molecules  Molecules are attracted to one another by a variety of forces.  These intermolecular forces are weaker than ionic or covalent bonds.  These forces are responsible for whether or not a molecular compound is a solid, liquid, or a gas.

45.  van der Waals forces – consist of dispersion forces and dipole interactions (dipole-dipole moments).  Dispersion forces – weakest of all intermolecular forces. They are caused by the motion of electrons. The strength of dispersion forces increases with the increasing number of electrons in a molecule.

46.  All molecules contain dispersion forces.  As molar mass and the number of electrons increase, dispersion forces increase.  Halogens are the most common molecules to have dispersion forces. Fluorine is a gas, Bromine is a liquid and Iodine is a solid.

47.  Dipole interactions – occur when polar molecules or ions are attracted to one another. This occurs when a partial positive charge and a partial negative charge come close to each other.  Dipole interactions are very similar to, but much weaker than ionic bonds.

48.  Hydrogen bonds – force exerted between a hydrogen atom bonded to an F, O, or N atom in one molecule and an unshared pair on another F, O, or N atom in a nearby molecule.  Hydrogen bonds can have a great effect on the boiling point of a substance.