Name Period Date Chemical Bonding & Intermolecular Forces (Chapter 12, 13 &14) Fill-in the blanks during the PowerPoint presentation in class. Ch. 12 Section 1: Introduction to Chemical Bonding Chemical Bonding are the electrons in the outer shell ( energy level) of an atom. A is a mutual electrical attraction between the and valence electrons of different atoms that binds the atoms together. During bonding, valence electrons are in ways that make the atoms more. The Three Major Types of Chemical Bonding results from the electrical attraction between oppositely-charged ions. results from the sharing of electron pairs between two atoms. results from the attraction between metal atoms and the surrounding sea of electrons. Ionic or Covalent? Bonding is usually somewhere between ionic and covalent, depending on the difference between the two atoms. In covalent bonds, the bonded atoms have an attraction for the shared electrons. Sample Question: Using electronegativity values (in table on pg. 362) to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs, and chlorine, Cl. In each pair, which atom will be more negative? Section 2: Covalent Bonding and Molecular Compounds Molecules A covalent bond is formed from of electrons. A is a neutral group of atoms held together by bonds.
Why Do Covalent Bonds Form? When two atoms form a covalent bond, their shared electrons form orbitals. This gives both atoms a stable configuration. The Octet Rule Atoms are the most when they have completely valence shells (like noble gases). The Octet Rule- Compounds tend to form so that each atom has an octet (group of ) electrons in it its highest energy level. is an exception to the octet rule since it can only have electrons in its valence shell. Electron-Dot Notation Electron-dot notation is indicated by dots placed around the element s. Only the electrons are shown. Inner-shell electrons are not shown. How to determine number of valence electrons : (On the periodic table) The number in front of A indicates the number of valence electrons for that group of elements. For example : Mg (group 2A) has valence electrons C (group 4A) has valence electrons Cl (group 7A) has valence electrons Xe (group 8A) has valence electrons Sample Problem: a) Write the electron-dot notation for hydrogen. b) Write the electron-dot notation for phosphorus. Lewis Structures Electron-dot notations of two or more atoms can be combined to represent. electrons will pair up to form a pair or covalent bond. The pair of dots representing the shared pair of electrons in a covalent bond is often replaced by a long. An unshared pair, also called a, is a pair of electrons that is not involved in bonding and that belongs to one atom.
How to Draw Lewis Structures 1. Draw the electron-dot notation for each type of atom, and count the valence electrons. 2. Put the least electronegative atom in the center (except H.) 3. Use electron pairs to form bonds between all atoms. 4. Make sure all atoms (except H) have octets. 5. Count the total electrons in your Lewis structure. Does it match the number you counted in step 1? If not, introduce multiple bonds. Sample Problem A: Draw the Lewis structure of iodomethane, CH 3 I. Multiple Covalent Bonds In a single covalent bond, pair of electrons is shared between two atoms. A bond is a covalent bond in which two pairs of electrons are shared between two atoms. A bond is a covalent bond in which three pairs of electrons are shared between two atoms. Multiple bonds are often found in molecules containing. Sample Problem B: Draw a Lewis structure for methanol, CH 2 O. Section 3: Ionic Bonding and Ionic Compounds Formation of Ionic Compounds Sodium and other metals easily electrons to form positively charged ions called. Chlorine and other non-metals easily electrons to form negatively-charged ions called. Ionic Bonding Cations (+) and anions (-) are attracted to each other because of their electrical charges. An ionic bond is a bond that forms between oppositely-charged ions because of their
electrical attraction. Ionic Bonding and the Crystal Lattice In an ionic crystal, ions minimize their energy by combining in an orderly arrangements known as a crystal. A is the smallest repeating unit of an ionic compound. Comparing Ionic and Covalent Compounds Covalent compounds have relatively forces of attraction between molecules, but ionic compounds have a attraction between ions. This causes some differences in their properties: Ionic Covalent Polyatomic Ions A charged group of covalently bonded atoms is known as a ion. Draw a Lewis structure for a polyatomic ion with brackets around it and the in the upper right hand corner. Section 4: The Metallic Bond In metals, overlapping orbitals allow the outer electrons of the atoms to throughout the entire metal. These mobile electrons form a around the metal atoms, which are packed together in a crystal. A results from the attraction between metal atoms and the surrounding sea of electrons. Properties of Metals The characteristics of metallic bonding gives metals their unique properties, listed below electrical thermal ( ) conductivity (can be hammered into thin sheets) (can be pulled or extruded into wires)
(shiny appearance) Section 5: Molecular Geometry VSEPR Theory The abbreviation VSEPR (say it Ves=pur ) stands for. VSEPR theory- repulsion between pairs of valence electrons around an atom causes the electron pairs to be oriented as as possible. Treat double and triple bonds as single bond. VSEPR theory can also account for the of molecules with electron pairs. VSEPR theory postulates that occupy space around the central atom just like bonding pairs, but they repel other electron pairs strongly than bonding pairs do. 2 electron pairs around a central atom will be 180 apart, and the molecule s shape will be. 3 bonding pairs around a central atom will be 120 apart, and the molecule s shape will be. If one of the pairs is a lone pair, the shape will be. 4 bonding pairs around a central atom will be 109.5 apart, and the molecule s shape will be. If one of the pairs is a lone pair, the shape will be. If two of the pairs are lone pairs, the shape will be. Sample Problem A: Use VSEPR theory to predict the molecular geometry of water, H 2 O. Unshared pairs repel electrons more strongly and will result in bond angles.
Sample Problem B: Use VSEPR theory to predict the molecular geometry of carbon dioxide, CO 2. Molecular Polarity Molecular polarity depends on both and molecular. o If all bonds are non-polar, the molecule is always non-polar. o If bonds are polar, but there is in the molecule so that the polarity of the bonds cancels out, then the molecule is (Ex: CO 2, CCl 4 ) o If bonds are polar but there is no symmetry such that they cancel each other out, the overall molecule is. (Ex: H 2 O, CH 3 Cl) Section 6: Intermolecular Forces The forces of attraction between molecules are called forces. Intermolecular forces vary in strength but are generally than any of the three types of chemical bonds (covalent, ionic or ). The strongest intermolecular forces exist between molecules. Because of their uneven charge distribution, polar molecules have. A dipole is represented by an arrow with its head pointing toward the pole and a crossed tail at the pole. Types of intermolecular forces (strongest to weakest): 1. - between 2 polar molecules. This (-) side of one dipole attracts the (+) side of another. a. bonding- a very strong type of dipole-dipole force. Only exists between atoms of H and. 2. - between a polar and a non-polar molecule. 3. dispersion force- instantaneous dipoles created by the constant motion of electrons. Ch. 13/14 State of Matter Section 1: The Kinetic-Molecular Theory The kinetic-molecular theory states: Particle of matter (atoms and molecules) are always in. We measure this energy of motion (kinetic energy) as. If temperature increases, the particles will gain more and move even.
Molecular motion is greatest in, less in, and least in. Gases An is a hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory. Many gases behave nearly ideally if is not very high and is not very low. - Gas particles glide easily past one another. Because liquids and gases, they are both referred to as. - Gas particles are very far apart. The density of a gas is about 1/1000 the density of the same substance in the liquid or solid state. - A gas will expand to fill its container. - The volume of a gas can be greatly decreased by pushing the particles closer together. Section 2: Liquids - strong forces at a liquid s surface act to decrease the surface to the smallest possible size. The higher force of between the particles of a liquid, the higher the surface tension. - A liquid or solid changing to a gas - particle escape from the of a liquid and become a gas. This occurs because liquid particles have kinetic energies. - bubbles of vapor appear liquid. Will not occur below a certain temperature (the point.) A liquid is on the evaporates readily. Section 3: Solids There are two main types of solids: Crystalline Solids- Made up of. Particles are arranged in an orderly, geometric, repeating. Solids- particles are arranged randomly. Melting Point- The temperature at which a solid becomes a. At this temperature, the kinetic energies of the particles within the solid overcome the forces holding them together.