Atoms have the ability to do two things in order to become isoelectronic with a Noble Gas.

CHEMICAL BONDING Atoms have the ability to do two things in order to become isoelectronic with a Noble Gas. 1.Electrons can be from one atom to another forming. Positive ions (cations) are formed when electrons are by metals. Negative ions (anions) are formed when electrons are by non-metals. The electrostatic attraction between the oppositely charged ions forms an bond. 2. Electron pairs can also be between two atoms. The mutual attraction of the electrons by each atom is called a bond. Valence electrons (outer electrons in the outermost energy level) of an atom are the electrons which in bonding. The number of an atom has determines how many other atoms it can bond with. Atoms with 1 4 valence electrons can form as many bonds as there are electron, eg. Carbon can from up to 4 bonds Atoms with 5 8 valence electrons can form up to [8 No. of valence] bonds, eg. Sulfur can from 8-6 = 2 bonds FORMING IONS A neutral atom can either gain or lose electrons to form an ion a charged species. Cations are named without changing the name of the element Example: A + Energy A + + e - Anions are named by removing the ending from the element and adding ide. Example A + e - A - + Energy Atoms can form many different ions but stable ions are those, which are (have the same electron configuration) with Noble Gas Ion Na - Ca 2+ F + P 3- P 5+ O 2+ Stable or unstable Unstable Stable Stable ion Name of Stable Isoelectronic ion with Na + Sodium ion Ne Ca 2+ Calcium ion Ar H1

IONIC BONDING Ionic bonding occurs when a combines with a. Metals lose electrons while non-metals gain electrons in order to form noble gas electron configuration. All the electrons lost by the metal are gained by the non-metal. Lewis (Electron Dot) Diagrams are used to show ionic bonding. Examples: Mg + O Li + O Al + Cl Formula unit: the simplest collection of atoms from which an ionic compound s formula can be established. A formula unit is to an ionic compound as a molecule is to a covalent compound, eg. NaCl or K 2 CO 3 Show the bonding between: Mg and P Write the formula unit for the salt. PROPERTIES OF IONIC SUBSTANCES Crystalline Ions are arranged in a crystal in a way to maximize the forces and minimize the forces between the ions. at room temperature melting points and boiling points but in water when dissolved in water [conduct electricity] As solids are of electricity As Liquids electricity H2

COVALENT BONDING- SHARING ELECTRONS Some elements do not lose or gain electrons easily. For these elements a different type of bonding occurs; bonding. Electrons are shared in pairs instead of being transferred. The basis of the bond is still attraction with electrons being attracted by two nuclei. Covalent bonds are formed when a combines with another. Since both nonmetals have to gain electrons, they achieve this by electrons. No are formed! The non-metals share electrons such that they have outer electrons when counting their own plus those "borrowed" from the atoms bonding with them. They mutually share electrons. Molecule Lewis diagram Molecule Lewis diagram H 2 OH - NH 3 CCl 4 H 2 O HCl PROPERTIES OF COVALENT SUBSTANCES Tend to form small groups of atoms called Exist as, or at room temperature Solids are often and Liquids and solids often are and can evaporate readily [ie. Perfume or mothballs] melting and boiling points solubility in water and other solvents Soluble in solvents such as chloroform or ether As solids, liquids or solutions are of electricity H3

MULTIPLE BONDS Up to electrons can be shared between atoms to achieve a stable octet in the valence orbital. Remember an atom can only share the number of valence electrons it has. Write Lewis Dot Diagram for: O 2 N 2 C 2 H 2 For more examples please read P. 64-67- Representing Molecular Structures Need & Have Method Need & Have Method- the have/need method will help you determine the number of bonds present in a molecule Show the Lewis dot diagram for the following: NH 4 + (aq) 1. Find out how many valence electrons are present: Have: 2. Find out how many valence electrons are needed to have a stable molecule: Need: 3. Take the difference of the two numbers and divide by 2 to determine the number of bonds: Difference = # of Bonds = 4. Draw the Lewis Dot Diagram: H4

COORDINATE COVALENT BONDS Co-ordinate covalent bonds are those, which bond when atoms share pairs of electrons One of the atoms provides pairs of electrons being shared while the other atom contributes electrons to the shared pair Example: Show the Lewis dot diagram for SO 3. Have: - Need: Difference = # of Bonds = RESONANCE STRUCTURES The of the position of the double bond can when the central atom is bonded to identical atoms This allows the double bond to and the net result is that each bond has the bond strength of. Electrons are delocalized Example: SO 3. H5

POLYATOMIC IONS 1. Polyatomic Ions always follow the octet rule. 2. Place the central atom with valance e -. 3. Place extra e - around the central atom. 4. Place satellite atoms with valance e -. 5. Do need/ have calculations. Put square brackets and net charge 2- Show the Lewis dot diagram for SO 4 (aq) Have: Need: Difference = # of Bonds = NON-OCTET rule for molecular compounds Many molecular compounds form, which do not follow the simple octet rule. The non- octet rule is applied when the atom (the element with the electronegativity) has or than eight electrons. In each such case, you will find that it is only the central atom (P or B below) that violates the octet rule. The surrounding atoms follow the octet rule (8e - ). There will be no coordinate covalent or multiple bonds. The central atom will have an even number of electrons around it. Example: PCl 5 BF 3 How to recognize Molecular Compound that does not follow the octet rule? Involves: 1. Boron (B); 2. Noble Gas, eg. XeF 4 ; 3. Two Halogens in a ratio greater than 1:1, eg. FCl 7 ; 4. Four or more atoms of one kind, eg. PBr 5 Additional Example: Show the Lewis dot diagram for compound that forms between element with 5 valence electrons and elements with 7 valence electrons. H6

POLAR COVALENT BONDS Ionic compounds form repeating. Covalent compounds form distinct. Consider adding to NaCl(s) vs. H 2 O(s): Sometimes when atoms of two different elements form a bond by sharing an electron pair there is sharing of electrons. When this occurs the bond is called a BOND. The unequal sharing results from the difference in of the two atoms. The one with the electronegativity exerts a attraction for the electrons ELECTRONEGATIVITY Recall that electronegativity is a number that describes the relative ability of an atom, when bonded, to Attract electrons. The periodic table has electronegativity values. We can determine the nature of a bond based on Δ EN (electronegativity difference). Δ EN = higher EN lower EN Example: NBr 3 Basically: a ΔEN below 0.4 = covalent, 0.4-1.7 = polar covalent, above 1.7 = ionic Determine the ΔEN and bond type for these: HCl, CrO, Br 2, H 2 O, CH 4, KCl H7

INTRAMOLECULAR FORCES Forces of electrostatic attraction a molecule. Occur between the of the atoms and their making up the molecule (i.e.covalent bonds). Must be broken by means. Form substances when broken INTERMOLECULAR FORCES Forces of attraction between molecules (i.e. London dispersion forces, dipole dipole Interactions or hydrogen bonds).these forces are much than Intramolecular forces or bonds and are much to break. Physical changes ( changes of ) break or weaken these forces Do form new substances when broken. These forces affect the and points of substances, the action and tension, as well as the and of substances Types of Intermolecular forces London Dispersion Forces London forces result from a type of dipole. These forces exist between molecules. They are masked by stronger forces (e.g. dipole-dipole) so are sometimes insignificant, but they are important in molecules. Because electrons are moving around in atoms there will be instants when the charge around an atom is not. The resulting tiny dipoles result in attractions between and/or. These forces are based on the simultaneous attraction of the of one molecule by the positive of neighbouring molecules The strength of the force is directly related to the of electrons and protons in a given molecule. The greater the number of electrons and protons the the force. Van der Waal force H8

Dipole - Dipole Interactions Occur between polar molecules having.molecules with dipoles are characterized by oppositely charged that are due to an distribution of charge on the molecule. The polarity of a molecule is determined by both the polarity of the bond and the of the molecule. These forces are based on the simultaneous attraction of the of one dipole by the of neighbouring molecules. The strength of the force is related to the of the given molecule Van der Waal force Hydrogen Bonds These forces are a of dipole dipole interaction. Occur between atoms in one molecule and electronegative atoms [F, O, and N] where there are usually unshared pairs of electrons present. Q- Calculate the EN for HCl and H 2 O The high EN of NH, OH, and HF bonds cause these to be strong forces. Also, because of the small size of hydrogen, it s positive charge can get very close to the negative dipole of another molecule. H-Bonding Diagram: Ionic Forces Ionic forces may be both inter and intra since a crystalline lattice is formed. For convenience sake ionic substances are referred to by the smallest ratio of atoms present in the lattice. H9

Predicting boiling points using the Strength of Intermolecular forces Molecules that are isoelectronic have the strength of London dispersion forces More polar molecules have dipole dipole interaction and melting and boiling points The the number of electrons per molecule, the the London forces and hence the the melting and boiling point Which would have a higher melting/boiling point? NaCl or HCl Question: For each, pick the one with the lower boiling point. Explain. a) CaCl 2, CaF 2 b) KCl, LiBr c) H 2 O, H 2 S. Polar Molecules Polar bonds cause the whole molecule to be polar. Polar molecule is a molecule in which the distribution of electrons results in a positive charge at end and a negative charge at the end. Non-polar molecule is a molecule in which the electrons are distributed among the atoms, resulting in no localized charges Bond dipoles may or may not cancel out thereby producing either molecules that are, if they, or, if they cancel. Example1: H10

Predicting Molecular Polarity General Steps Go over Tutorial 1 and Table 3 on pg. 106 and use the VSEPR THEORY to guide you. Step 1: Draw a reasonable Lewis structure for the substance. Step 2: Identify each bond as either polar or nonpolar. (If the difference in electronegativity for the atoms in a bond is greater than 0.4, we consider the bond polar. If the difference in electronegativity is less than 0.4, the bond is essentially nonpolar.) If there are no polar bonds, the molecule is nonpolar. If the molecule has polar bonds, move on to Step #3. Step 3: If there is only one central atom, examine the electron groups around it. If there are no lone pairs on the central atom, and if all the bonds to the central atom are the same, the molecule is nonpolar. If the central atom has at least one polar bond and if the groups bonded to the central atom are not all identical, the molecule is probably polar. Move on to Step #4. Step 4: Draw a geometric sketch of the molecule. Step 5: Determine the symmetry of the molecule using the following steps. Describe the polar bonds with arrows pointing toward the more electronegative element. Use the length of the arrow to show the relative polarities of the different bonds. (A greater difference in electronegativity suggests a more polar bond, which is described with a longer arrow.) Decide whether the arrangement of arrows is symmetrical or asymmetrical If the arrangement is symmetrical and the arrows are of equal length, the molecule is nonpolar. If the arrows are of different lengths, and if they do not balance each other, the molecule is polar. If the arrangement is asymmetrical, the molecule is polar. Predicting Molecular Polarity- Exercises Decide whether the molecules represented by the following formulas are polar or nonpolar. (You may need to draw Lewis structures and geometric sketches to do so.) a. CO 2 b. OF 2 c. CCl 4 d. CH 2 Cl 2 e. HCN H11

VSEPR Theory - Valence Shell Electron Pair Repulsion Theory General Formulas Every molecule can be described using a general formula where: A = Central Atom B = Bonding electron pairs (surrounding atoms) E = non-bonding lone electron pairs Every general formula results in the same 3D structure: General Formula Structure Name Example AB 2 E 0 or AB 2 Linear CO 2 AB 3 E 0 or AB 3 Trigonal Planar BF 3 AB 4 E 0 or AB 4 Tetrahedral CH 4 AB 5 E 0 or AB 5 Trigonal Bipyramidal PF 5 AB 6 E 0 or AB 6 Octahedral XeF 6 Five main categories of structures exist termed families. In each family, as the number of lone electron pairs (E) increases, the shape of the molecule changes slightly. Tetrahedral Family General Structure Name Example Formula AB 4 E 0 Tetrahedral CH 4 AB 3 E 1 Trigonal NH 3 Pyramidal AB 2 E 2 Bent/V-Shaped H 2 O AB 1 E 3 Linear HCl General Formula AB 5 E 0 Trigonal Bipyramidal Family Structure Name Trigonal bipyramidal Example PF 5 AB 4 E 1 See-saw SF 4 AB 3 E 2 T-shaped ICl 3 AB 2 E 3 Linear XeF 2 Octahedral Family General Formula Structure Name Example AB 6 E 0 Octahedral SF 6 AB 5 E 1 Square pyramidal BrF 5 AB 4 E 2 Square planar XeF 4 Steps to Determining the Shape of a Molecule: 1) Draw the Lewis structure of the molecule. 2) Analyze the Lewis structure to determine the general formula (A, B, E) 3) Match the general formula to the shape name 4) Determine the molecular polarity H12