Chemical Bonding PDF

IONIC BONDING Chemical Bonds are links of atoms that form different kinds of matter \* Two kinds of chemical bonds are: 1\. lonic Bonds: chemical bond resulting from the transfer of electrons from one bonding atom to another. - Static electrical attraction (electrostatic) is the basis for...

IONIC BONDING Chemical Bonds are links of atoms that form different kinds of matter \* Two kinds of chemical bonds are: 1\. lonic Bonds: chemical bond resulting from the transfer of electrons from one bonding atom to another. - Static electrical attraction (electrostatic) is the basis for ionic compounds - In an ionic bond, a positively charged ion (cation) is attracted to a negatively charged ion (anion). - lonic Compounds: compound of + and - ions combined so that the charges are neutralized; formed from a metal and a nonmetal. (polyatomic ions also) 2\. Covalent Bonds: chemical bond resulting from the sharing of electrons between two bonding atoms. PROPERTIES OF COVALENT BONDS - Polar: Description of a bond that has an uneven distribution of charge due to an unequal sharing of bonding electrons. - Nonpolar: Description of a bond that has an even distribution of charge due to an equal sharing of bonding electrons. - Electrons are shared unequally in a polar covalent bond and equally in a nonpolar covalent bond. The bonding atoms difference in electronegativity indicates whether the bond is ionic, polar, or nopolar. - Electronegativity difference determines the polarity of a bond. - \*The polarity of a bond is determined by the electronegativity difference between the 2 atoms. COVALENT BONDING - A Covalent Bond is formed by a shared pair of electrons between two atoms. - Molecule: A group of atoms that are united by covalent bonds. - Molecular Substance: Is a substance made of molecules. Molecular Formula: Tells how many atoms are in a single molecule of the compound. This is used to describe the composition of a molecular compound. Bond Energy - the energy needed required to break a chemical bond and for neutral isolated atoms. THE OCTET RULE \* States that atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons, which is 8, seen as s? po The ions in an ionic compound and the atoms in a molecular compounds satisfy the Octet Rule. EXCEPTIONS TO THE OCTET RULE 1\. Atoms with less than an octet. - H and He happy (stable) with 2 valence electrons→ makes sense since they have the 1s only. - Be happy (stable) with 4 valence electrons→ makes sense since he has the 1s and 2s. - B is the weirdo he is happy with 6 valence electrons. 1. Atoms with more than an octet (some atoms found beyond the second row of the periodic table form bonds that give them more than an octet of electrons). Total can be 10 or 12 valence electrons. Only do this when they have to. Must have d sublevel to do this. It is called an expanded octet. This is what happens when you have too many electrons, the extra go to the atom in the middle. 2. Molecules with an odd number of electrons. You can\'t make an octet with an odd number of electrons. LEWIS DOT DIAGRAMS - In Lewis dot diagrams, dots are placed around an element symbol to represent an atom\'s valence electrons. - The central atoms should be the one in the front, the lower electronegativity atom, and the one that gives the molecule symmetry. - At least 2 of 3 should agree for the above statement. Lewis Structures: Type of structural formula that uses dots or dashes to indicate bonds. - Unshared Pair: Pair of electrons that is not involved in bonding but instead is held exclusively by one atom. - Atoms in a molecule may satisfy the octet rule by forming single, double, or triple covalent bonds. Single covalent bonds: Chemical bond resulting from the sharing of an electron pair between two atoms. MULTIPLE BONDS- this is what happens when you run out of electrons before everyone has an octet or is stable. \* Double covalent bonds: Consists of two pairs of shared electrons= 4 electrons. This occurs when you are 2 electrons short of everyone having an octet. \*Triple covalent bonds: Three pairs of shared electrons. = 6 electrons. This occurs when you are 4 electrons short of everyone having an octet. Resonance- refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure. Occurs when a multiple bond can be correctly drawn in more than one location. TYPES OF IONS Monatomic lons: One atom ion; An ion formed by one single atom. Polyatomic lons: An ion made from more than one atom, charged group of covalently bonded atoms. BINARY IONIC COMPOUNDS Binary lonic Compounds: lonic compound consisting of two elements. Empirical Formula: Chemical formula that gives the simplest whole-number ratio of atoms of elements in a compound. \* An empirical Formula indicates the relative \# of atoms or ions in a compound. - For binary ionic compounds, the empirical formula can be written using the lowest common multiple. - Formula unit- smallest piece (representative particle) of an ionic compound that still retains the properties of that compound. Metallic bonding- chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons. Malleability- to be able to be hammered into thin sheets. Ductility- to be able to be drawn into wires. VSEPR- (valence shell electron pair repulsion) states that the repulsion between the sets of valence- level electrons surrounding an atom causes these sets to be oriented as far apart as possible. This is done to minimize repulsion. This causes the various shapes and bond angles to occur. Hybridization- the mixture of two or more atomic orbitals of similar energies on the same atom to produce new hybrid orbitals of equal energies Hybrid orbitals- orbitals of equal energy produced by the combination of two or more orbitals on the same atom. Hybrid Orbital- orbital of electrons in a bond, which is a combination of the shapes and the properties of the original orbitals in the atom. - sp Orbital- combination of one s-orbital and one p-orbital - sp? Orbital- one s- and 2-p-orbitals - sp3 Orbital- one s- and 3-p-orbitals\ Hybrid orbitals have properties of both kinds of orbitals. Atoms farther down a group of the periodic table form longer bonds. Multiple bonds are Shorter than single bonds, because bonds with more electrons attract the nuclei of the bonding atoms more strongly Repulsion Ranking- how electrons repel when in pairs. - 2 Unshared pairs- greatest type of repulsion - 1 Shared, 1 Unshared- second greatest - 2 Shared pairs -weakest type of repulsion → so therefore the shapes on the VSEPR charts are due to the species trying to minimize repulsion. Bond Angle- a geometric angle between two adjacent bonds. dipole- created by equal but opposite charges that are separated by a short distance. Counting the shared: unshared ratio of electron domains. 1. center 2. bigger atom 3. mirror image Counting the shared: unshared ratio of electron domains. - single, double, or triple bonds, all of them count only once for the shared: unshared ratio. You label it one adjective or the other. - when drawing the dot diagram if you run out of electrons before they all are happy (octet) you will use a multiple bond. If you have too many electrons you will get an expanded octet. - the atom with the lower electronegativity usually goes in the middle, or the atom in front of the formula, or the one that gives molecule symmetry. Most of the time it all agrees. Recall from before\.... - The polarity of a bond is determined by the electronegativity difference between the 2 atoms. - \*\*they use the same term polarity to describe a molecule. - The polarity of a molecule is determined by two things: The polarity of its bonds, and its shape. - When two different polarity directions exist, they cancel each other out if the force is equal, and then the molecule is nonpolar. Intermolecular forces- vary in strength but are generically weaker than bonds that join atoms in molecules, ions in ionic compounds or metal atoms in solid metals There are 3 kinds of intermolecular forces. They are, in order of increasing strength: 1\) London dispersion forces- the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles. Criteria-all molecular compounds and noble gases have LDF. It is considered a temporary or induced dipole. The molecule literally becomes distorted. LDF is directly proportional to size (number of electrons) = tie breaker 2\) Dipole- dipole forces= /interaction dipole- is created by equal but opposite charges that are separated by a short distance Dipole: If a molecule has a positive end and a negative end, it is also called a polar molecule; nondipoles are nopolar. Criteria- must be a polar molecule. tie breaker= electronegativity 3\) Hydrogen bonding- the intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom (NOF) is attracted to an unshared pair of electrons of NOF in a nearby molecule. NOF= nitrogen, oxygen, or fluorine. tie breaker= electronegativity \*if the intermolecular force of attraction is stronger, the following properties will also increase. Boiling point, melting point, surface tension, capillary action, and viscosity. 1\. Define chemical bond\ 2. Which electrons bond?\ 3. Bonding causes atoms to lower their potential energy and become more stable\ 4. Polar vs nonpolar bond (2)\ 5. Define electronegativity\ 6. Covalent vs ionic bond (2)\ 7. Electronegativity difference for each type of bond\ 8. Assign partial charge\ 9. Define molecule\ 10. Define molecular formula\ 11. Define bond length\ 12. Use octet rule\ 13. What has an octet already?\ 14. Who goes in the center of a dot diagram?\ 15. Define resonance\ 16. Ions in an ionic compound make what shape?\ 17. Define lattice energy\ 18. Higher lattice energy means stronger \_\_\_\_\_\_\_\_?\ 19. Ionic vs molecular compounds (2)\ 20. Define metallic bonding\ 21. Properties of metals\ 22. VSEPR is used for?\ Free response. Do the big six for the following.\ 1. Dot diagram 2. Shared:unshared 3. Shape\ 4. Bond angle 5. Hybridization 6. Polarity of molecule\ a. H2S b. NF3 c. CI4 d. HCN e. OF2 f. PBr3 g. Cl2 h. CH2O i. BrCl j. CO2