Chemistry Chapter 6 A

Describe the role of each of the following in predicting molecular geometries: unshared electron pairs and double bonds.

Unshared electron pairs occupy space as bonded electrons, but they are not part of the visualized by molecular geometry. Double bonds are treated the same and single.

What happens to a liquid's kinetic energy and boiling point? What is the relationship between boiling point and molecular forces?

a. The KE increases and at boiling point, the energy is above to overcome the force of attraction between the liquid's particles. b. The higher the boiling point, the stronger the forces between the molecules.

Name the weakest to strongest. Covalent, ionic, dipole dipole, London dispersion

4, 3, 2, 1

How are dipole-dipole attractions, London dispersion forces, and hydrogen bonding similar?

In all cases there is an attraction b/w slightly negative portion of one molecule and slightly positive portion of another.

In what forms do such ions often occur in nature?

Polyatomic ions combine with ions of opposite charge to make ionic compounds.

London Dispersion

Act between all atoms and molecules but they are the only intermolecular forces acting among noble gas atoms and nonpolar molecules. This is reflected in the low boiling points of NG and NP molecules. London is dependent on the motion of electrons, their strewth increases with the number of electrons meaning it increases with increasing atomic or molar mass.

What types of atoms tend to form the following types of bonding? Ionic Covalent Metallic

Lattice, metals and non metals Bond, nonmetals only Heat of Vaporization, metals only

Components, overall charge, conductivity, melting point, hardness, malleable, ductile of metallic, ionic, and covenant bonds.

Metallic: Atoms, neutral, yes, L-h, soft-hard, yes, yes Ionic: Ions, -neutral, yes, high, hard, no, no. Covalent: Atoms, neutral, no, low melting points, n/a, no

1. Best way to describe structure of solid metals? 2. Why are electrons are free to move throughout a piece of metal? 3. What does bond strength determine? 4. The heat of vaporization of aluminum is 294 kJ/mol, and the heat of vaporization of sodium is 97 kJ/mol, about 1/3 that of aluminum. From these data you can predict that: and why? 5. What does the chemical formula in ionic compounds show? 6. What happens to 2 different elements as they form an ionic compound?

1. Atoms arranged in a crystal lattice surrounded by electrons that are free to move. 2. Because some of the outer electrons of the metal atoms are delocalized. 3. Hardness 4. The bond strength in aluminum is greater than the bond strength in sodium is. Heat of vaporization is an indicator of bond strength in metals, so the bond strength of aluminum is greater than that of sodium. 5. The ratio of positive and negative ions in a sample of the compound. 6. One element loses electrons to atoms of the other element. Atoms of the element with lower electronegativity lose electrons to the element with higher electronegativity, and, usually, the ions formed have noble gas configurations.

1. What are the types of chemical bonds? 2. What is a chemical bond? 3. Ionic bonding and Covalent bonding? 4. Non-polar covalent bond? 5. Polar-covalent bond?

1. Covalent, covalent network, ionic, metallic 2. A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. Most atoms are at a lower energy state when bonded to other atoms rather than as independent particles. 3. Ionic bonding results from the electrical attraction between large numbers of cations on anions. Electrons are transferred. Covalent bonding results from the sharing of electron pairs between 2 atoms. 4. Bonding electrons are shared equally by bonded atoms, resulting in a balanced distribution of electrial charge. 5. Bonding electrons are shared unequally because one atom pulls although not very strong, more on the electron pair than the other atom.

1. Electron dot diagrams? 2. Lewis structures? 3. Structural formula? 4. Resonance structures 5. Covalent-network bonding

1. Electron configuration notation in which only the valence electrons of an atom of a particular element are showed, indicated by dots. 2. Formulas in which atomic symbols represent nuclei and inner shell electrons. Dot pairs or dashed between 2 atoms represent electron pairs in covalent bond. Dots adjacent to only one atomic symbol are unshared electrons. 3. Indicates the kind, number, arrangement, and bonds but no the unshared pairs of the atoms in the molecule. 4. Bonding in molecules or ions that cannot be corrected represented by a single Lewis structure. 5. Covalent bonding exists throughout a large network of atoms.

1. What are some other properties of ions? 2. What is the VSEPR Theory? 3. Repulsion between electron pairs? 4. Hybridization Theory 5. Intermolecular force. 6. 3 kinds of intermolecular forces +functions 7. Polar and non-polar molecules in symmetry. 8. Hydrogen bonding?

1. Hardness: tough for one layer to slide past another, Brittle: if ionic layers do shift, they snap. 2. Repulsion between the sets of valence level electrons surrounding an atom causes these sets to be oriented as far apart as possible. Stands for valence shell, electron pair and repulsion. 3. Unshared pairx2 > unshared-shared > shared-shared. Unshared occupies more space and shared less. 4. Explains how the orbitals of an atom rearrange when covalent bonding forms. Mixing of 2 or more orbitals of similar energies on the same atom to produce new orbitals of equal energies. 5. The forces of attraction between molecules. Generally weaker than other forms of bonding. Boiling point value is a good measure of intermolecular attraction because as B. PT gets higher, forces between particles get higher. 6. Dipole is created by equal but opposite charges that are separated by a short distance. For example, ---> The 3 are dipole dipole, induced dipole and London dispersion. Dipole dipole: force of attraction between polar molecules. Induced is nonpolar molecule is changed into a dipole when its electron cloud is distorted by an approaching dipole. London dispersion forces: temp dipoles form when electrons, constantly in motion, momentarily shrift unevenly. Act between all atoms and molecules and increase with the # electrons. 7. Polar bonds can produce both polar and non polar molecules. Analyze geometric shape to determine if polarities cancel out. Polar molecules with unequal charge distribution and non polar with equal distribution. 8. Specific example of a strong dipole dipole bond. Isn't really a chemical bond, drawn with dotted lines.

1. How does the behavior of electron in metals contribute to the metal's ability to conduct electricity and heat? 2. What is the relationship between the heat of vaporization of a metal and the strength of the bonds that hold the metal together?

1. Mobility o/the electron in a network of metal atoms contributes to the metal's ability to conduct electricity+ heat. 2. Amount of heat required to vaporize metal is a measure of the strength of the bonds that hold the metal together.

1. When do atoms have high electronegativity? 2. Ionic character relationship? 3. How can electronegativity be used to distinguish between an ionic bond and a covalent bond? 4. What best represents the strength in ionic bonds? 5. What are valence electrons attached to in metallic bonds? 6. When light strikes the surface of a metal, what happens? 7. What are mobile electrons in metallic bonds responsible for? 8. What happens to the strength of the metallic bond when moving left to right on the periodic table? 9. What happens when metal is drawn into wire? 10. Use the concept of electron configuration to explain why the number of valence electrons in metals tends to be less than the number in most nonmetals.

1. When they have a strong attraction for electrons they share with another atom. 2. The greater the electronegativity difference between two atoms bonded together, the greater the bonds percentage of ionic character. 3. The larger the difference in electronegitivity, the more ionic. The smaller, the more covalent. If the difference is less than 1.7, it's covalent. If it's more, it's ionic. 4. Lattice Energy. 5. Shared by surrounding atoms. 6. The electrons in the electron sea absorb and re-emit the light. 7. Luster, thermal conductivity, and electrical conductivity. 8. It increases. 9. Metallic bonds do not break. 10. Most metals have their other electrons in s orbitals non metals in p. The more localized the electrons, the harder the metal.

6. Define attraction, repulsion, and effect of both. 7. Molecule, molecular compound, molecular formula, and diatomic molecule. 8. Bond length? 9. Bond energy? How is it measured? 10. Octet rule?

6. Attraction the nucleus of one atom to the electron. Repulsion is when both nuclei repel, as well as both electron clouds. As the atoms approach, attraction increases and potential E goes down. Beyond a certain point, repulsion increases and R goes up. Therefore, bottom of valley on E curve is where there is a balance between attraction and repulsion. 7. Molecule is a neutral group of atoms help together by covalent bonds. A molecule compound is a compound that is made of molecules. Molecular formal shows the kinds and numbers of atoms making up a molecule. Diatomic molecule is a molecule containing only 2 atoms. 8. Bond length is the average distance bw 2 bonded atoms. 9. Energy required to break a chemical bond and form neutral isolated atoms. Measured in KJ/mole. It has a positive value. This is because bond energy is giving off energy and getting it back. 10. Chemical compounds tend to form so that each atom, by gaining, losing or sharing electrons, has an octet of electrons in its highest energy level. H/B exceptions.

6. Ionic compound and formations of them. 7. Simplest unions of ionic and covalent bonding? 8. Lattice Energy 9. Bond energies when new bonds make/broken? 10. Ionic vs. Covalent properties

6. Composed of positive and negative ions that are combined so that the number of positive charges are equal to the number of negative charges (cancel out). Called formula units. Formation: Atoms form ions when they lose/gain in order to attain a noble gas configuration. Ions reach a lower potential E through electrical forces of attraction between oppositely charge particles when they combine in an orderly arrangement known as a crystal lattice. Forces of repulsion (like charged ions, adjacent e-clouds) balance the forces of attraction (opp. charged ions, nuclei-electron of adjacent ions). 7. Molecule, formula unit. 8. Term used to describe bond strength in ionic compounds. Energy released when one of an ionic crystalline compound is formed from gaseous ions. Negative value indicated that E is released. The larger the value, the move E released, the more stable bonding will be. 9. Put energy to break bond, energy given off when a new bond is made. 10. Ionic: strong force that hold ions together, stronger forces of attraction between ions therefore higher melting points and boiling points, don't vaporize in room temperature. Covalent: Strong covalent bond within the atoms of each molecule, weaker F of attraction between molecules, therefore melt at low temps and many vaporize at room temperature.

What two factors determine whether or not a molecule is polar?

Electronegativity difference and molecular geometry.

Metallic bond strength What is bond energy, lattice, and heat of vaporization?

Expressed in the heat of vaporization where the bonded atoms in the metallic slid state are converted into individual metal atoms in the gaseous state (usually ^ heat of of vaporization, the ^ the bond strength) determined by strength of nuclear charge and # delocalized electrons Bond e- is the energy added to break a covalent bond Lattice e- is released when ionic compounds are broken into atoms Heat of vaporization e- added when bonded metallic solid atoms are broken into individual gas atoms.

Intermolecular Molecular polarity and dipole-dipole forces.

Forces of attraction between molecules. Weaker than bonds that join atoms in molecules, ions in ionic compounds, or metal atoms in solid metals. Ionic compounds and metals have higher boiling points than molecular substances. Strongest in polar molecules. They act as tiny dipoles because of their uneven charge distribution. A dipole is created by equal but opposite charges that are separated by a short distance. Dipoles are represented by arrows. The forces of attraction between polar molecules are know as dipole dipole forces. These are short ranged, acting only on nearby molecules. Polarity of diatomic molecules (BrF) are determined by just one bond. For molecules containing more than 2 atoms, molecular polarity depends on the polarity and orientation of each bond. A polar molecule can induce a dipole in a non polar molecule by temporary attracting its electrons. The result is a short rang intermolecular force that is somewhat weaker than the dipole dipole force.

Metallic Bond Model

Metals have very few electrons in their highest E level. They have many vacant d orbitals below the outer level. Vacant orbitals of adjacent atoms overlap which allows these loosely held e-s to roam freely. Delocalized electrons don't stay in one locality like covalent bonding (stay in the overlapping of shared orbitals) or ionic bonding where electrons are bound to an ion within crystal lattice. Mobile electrons form a sea of electrons.

Numbers: non polar, polar, ionic.

Nonpolar: 0-0.3 Polar covalent: 0.3-1.7 Ionic: 1.7+

H2S and H2O have similar structures and their central atoms belong to the same group. Yet H2S is a gas at room temperature and H2O is a liquid. Use bonding principles to explain why this is true.

O2 is more electronegative than sulfur which creates a more polar bond. Increased polarity in H2O bonds means a stronger intermolecular attraction, making it a liquid in room temperature.

In drawing the Lewis structure the central atom is generally the?

The central atom in a Lewis structure is usually the least electronegative atom.

Metallic bonding and properties Luster, malleability, ductility

The chemical bonding that results from the attraction between metal atoms and surrounding electron sea. Mutual sharing of many electrons where each atom contribute its valence electrons which are then free to move about the mostly vacant outer orbitals of all the metal atoms. High electrical and thermal conductivity due to high mobility and delocalization of electrons. Luster (shine) metals absorb energy and become "excited" very easily because of many of their orbitals are separated by extremely small E..shine occurs when photons are emitted when excited electrons return to the ground state. Malleability (ability to be hammered/beaten into thing sheets) Ductility (ability to be drawn, pulled, or extruded to produce wire) because metallic bonding is the same in all directions and a shift in layers of atoms is inconsequential.

Why is a polar-covalent bond similar to an ionic bond?

The difference of electronegative between the two atoms in both types of bonds results in electrons being more closely associated with the electronegative atom.

In general what determines whether atoms will form chemical bonds?

The electron arrangement of the outer energy level of an atom determines whether or not it will form chemical bonds.

In water, 2 H bonds are bonded to one O. Why isn't it linear?

The electron pairs that are not involved in bonding also take up space creating a tetrahedron of pairs.

Order from strongest to weakest attraction: a. Polar molecule and polar b. Nonpolar and nonpolar molecule c. polar and ion molecule d. ion and ion

d, c, a, b

Hydrogen bonding

The electronegative differences between hydrogen atoms and FON make the bonds connecting highly polar. This gives the hydrogen a positive charge that is almost half as large as that of a proton. The small shape of the hydrogen allows it to become close to unshared pairs of electrons an an adjacent molecule. The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.

What is the relationship between electronegativeity and the ionic character of a chemical bond?

The greater the electronegativity difference, the greater the percent ionic character in a bond.

What is relationship between lattice energy and the strength of ionic bonding?

The more lattice energy there is, the more the ionic bond attracts electrons from other atoms forming new compounds.

Describe the octet rule in terms of noble-gas configurations and potential energy?

The octet rule is a simple chemical rule of thumb that states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas.

How is VSEPR used to classify molecules?

The shapes of the molecules are classified based on the number of bonding electron pairs and lone pairs that surround a molecule's central atom.

Describe the general location of electrons in a covalent bond?

The shared electrons are typically near the middle of the bond between the 2 atoms, in a covalent bond. They may be slightly closer to 1 atom or the other, due to small differences in electronegativity.

What two theories can be used to predict molecular geometry? What are some factors that affect the geometry of a molecule? What type of intermolecular force contributes to high boiling point of water?

Valence shell electron pair repulsion (VSEPR) theory and Electron Domain (ED) Theory can predict the molecular geometry. Electron pairs in a molecule want to adopt a geometry to maximize their distance apart, since like charges repel each other. Hydrogen bonding in water keeps the molecules attracted to one another and raises the boiling point.

What determines the number of hybrid orbitals produced by an atom? How do intermolecular forces compare in strength with those in ionic and metallic bonding? Where are the strongest intermolecular forces found?

a. Always equal to the number of orbitals that have combined b. Intermolecular forces are weaker than the forces involved in ionic and metallic bonding. c. In polar molecules

What is hydrogen bonding? What accounts for its strength? What are London dispersion forces?

a. Hydrogen bonding is a particularly strong dipole dipole force that occurs among molecules containing hydrogen atoms and highly electronegative atoms like N, O, Cl, and F. The great electronegativity difference between H and F N O or Cl an atoms of hydrogen has a positive charge approaching that of a proton. This, coupled with the small size of the hydrogen atom, results in a very strong dipole dipole attraction. b. London dispersion forces are intermolecular forces resulting from the creating of instantaneous dipoles.

What is the relationship between electronegativity and the polarity of a chemical bond? What determines the polarity of a molecule? Why are induced dipoles important?

a. The more electronegative atom in a covalent bond draws electrons toward it, creating a polar bond. b. By the polity of the molecule's individual bonds as well as the orientation of the bonds with respect to one another. c. Induced dipoles account for the solubility of non polar compounds such as O2 in polar compounds such as H2O

What happens to the energy level and stability of two bonded atoms when they are separated and become individual atoms? How are ionic and covalent bonds different? How are ionic and molecular compounds different? How does an ion differ from a metal? How does the energy level of a hybrid orbital compare wit the energy levels of the orbital it was formed from?

a. They energies of the atoms increase and the atoms become less stable b. In ionic, valence electrons of the atoms of the less electronegative element are donated entirely to the atoms of the more electronegative element. In covalent, valence are shared between the bonded atoms. c. A molecular compound consists of individual units capable of existing on their own. An ionic compound consists of an arrangement of a large number of ions. There is no discrete, independent particle in an ionic compound. d. An ionic pound is held together by electrical attraction between ions. A metal is held by sharing of atoms of a sea of mobile valence electrons. e. It lies between the energy levels of the orbitals from which it was made.