Chem 20 - Bonding

Predicting Molecular Shapes

1. Draw out the lewis structure of the molecule based on the chemical formula. 2. Determine the total number of electron domains around the central atom. This will be based on the number of lone pairs, single bonds, double or triple bonds that are formed. 3. Draw a structural formula that represents one of the shapes covered. The repulsive forces of the electron pairs will push the peripheral atoms as far apart as possible from the central atom. Unbonded pairs can take up more room within an orbital and exert more repulsive forces on the peripheral atoms.

Drawing Lewis Structures of Coordinate Covalent Bonds

1. Draw the Lewis structure for the compound as usual, but also use a lone pair to make a bond. 2. If the compound is negatively charged, then there is a gain of an electron to complete the lone pair. Represent this on the dot diagram and place brackets around the compound and the charge on the outside. 3. If the compound is positively charged, then an electron is lost since only the nucleus of the element was attracted. Represent this on the dot diagram and place brackets around the compound and the charge on the outside.

Drawing a Lewis Formula (compound)

1. Draw the Lewis symbol of the element that has the most unfilled orbitals (highest bonding capacity). This is the central atom. 2. Add the Lewis symbol of other elements (peripheral atoms) in the compound by making electron pairs to fill the empty orbitals. The molecular formula may provide some hints on the arrangement. 3. Incomplete orbitals within the central atoms is often a sign of double or triple bonds. For double or triple bonds, draw the electron pairs within the central atom.

Drawing Lewis Structures of Ionic Compounds

1. Draw the metal cation in square brackets with the positive charge outside of the bracket. 2. Draw a Lewis dot diagram with completely filled valences for the non-metal anion in square brackets with the negative charge outside the square brackets. 3. Draw as many metal and non-metal atoms that exist in the compound.

Octet Rule

According to the octet rule, elements can only have a total of 8 valence electrons. When an element has a filled, outermost energy level, it is at its most stable. This explains why noble gases are highly unreactive. Exceptions to the octet rule include all elements from Hydrogen to Boron.

Properties of Alloys

Alloys tend to be stronger and stiffer than pure metals. This is because the different sized atoms (cations) prevent planes of metals sliding past each other as opposed to a pure metal. Alloys can also combine properties of pure metals which may or may not be advantageous.

Ionic Compounds

Are usually compounds that are composed of a metal and a non-metal. Exceptions to this are compounds composed of polyatomic cations and polyatomic anions (ionic bonding).

Molecular Compounds as Solids

As solids, molecular compounds can be very soft (eg. wax) as the attractions holding the molecules together are London Dispersion forces arising from induced dipoles, or can be hard and brittle (eg. sugar) due to permanent dipole-dipole interactions. The hardest molecular solids derive their strength from the network arrangement and contribution of covalent forces together (diamonds and silica - giant covalent structures)

Electronegativity and Size

As you move from left to right within any given period, more protons are added to the nucleus, which increases the attractive force and pulls the energy level closer to the nucleus. The greater attractive force (due to a more positively charged nucleus) and smaller radius (less distance) results in a higher E.N. The attractive force between opposite charges decreases with the square of the distance between the charges (2X distance: 4X less attraction) Because of the small radius of atoms with a high electronegativity, electrons from an adjacent atom can get closer to the nucleus of the smaller atom.

Dipole-Dipole Interactions (Permanent Dipoles)

Atoms with high electronegativity can attract electrons more strongly resulting in one end of a molecule being more negative and the other end being more positive (polar). The positive and negative ends are referred to as dipoles. The positive dipole of a molecule can interact with the negative dipole of a neighbouring molecule, while the negative dipole of that molecule can interact with the positive dipole of another neighbouring molecule. These interactions can orient the molecules within a molecular compound to form a regular structure that is similar to that of an ionic compound's crystal lattice. . Although these forces are not as strong as the charge attraction of ionic compounds, they can still stabilize polar molecules into a solid crystal. Eg. sucrose (table sugar)

Delocalized Electrons and Malleability

Because a sea of delocalized electrons holds the metal atoms together, there is no directional attraction within the solid. As such, when stress is applied to metals, the sea of delocalized electrons can shift to accommodate the sliding of the positive metal ions past each other, maintaining the active forces holding the solid together, and resulting in metals being malleable and ductile.

Electron Domain

Because the attractive forces behind the bond formation and the repulsive nature that influences a compound's shape reside within the pairing of electrons, the capacity for bonding and repulsion within a given orbital on the central atom within a compound can be referred to as an electron domain/charged centres.

Shape of Boron Trihydride (trigonal planar)

Because the bonding that occurs within this compound involves three orbitals, this compound contains three electron domains. The repulsive forces of the three electron domains will push the peripheral atoms as far away from each other as possible, resulting in three 120 degrees angles and a trigonal planar molecule. There are no unbonded pairs of electrons to influence the molecular shape as all orbitals are involved in bonding (boron is an exception to the octet rule) This can be represented by the symbol AX3, where A represents the central atom with no unbonded pairs and X represents any atom bonded to the central atom.

Shape of HCN (linear)

Because the bonding that occurs within this compound involves two orbitals on the Carbon (central atoms), this compound contains two electron domains. The repulsive forces of both electron domains will push the peripheral atoms as far away from each other as possible, resulting in a 180 degrees angle and a linear molecule. There are no unbonded (lone) pairs of electrons to influence the molecular shape and that is made possible by the formation of a triple bond. This can be represented by the symbol AX2, where A represents the central atom with no unbonded pairs and x represents any atom bonded to the central atom.

Electronegativity and Ionic/Covalent Bonds

Because there is always some attraction between the nucleus of one atom and the electrons of another, there is no clear distinction between ionic and covalent bonds. Can be viewed on a continuum relating mostly ionic character vs mostly covalent character.

Shape of Methane (tetrahedral)

Bonding involves all 4 orbitals, resulting in 4 electron domains. The repulsive forces of the 4 electron domains will push the peripheral atoms equally away from each other, resulting in a 109.5 angle and a tetrahedral molecule. Since all orbitals are involved in bonding, there are no lone pairs left over to influence the shape of the molecule. This can be represented by the symbol AX4, where A represents the central atom, and X represents any atom bonded to the central atom.

Shape of Ammonia (pyramidal)

Bonding involves three orbitals, but the repulsive forces in an unbonded pair results in 4 electron domains. The unbonded pair will exert more of a repulsive force, altering the shape (bending peripheral atoms) by creating a 107.3 degree angle and a pyramidal molecule. This can be represented by the symbol AX3E, where A represents the central atom, C represents any atom bonded to the central atom, and E represents an unbonded pair..

Shape of Water (bent - tetrahedral group)

Bonding involves two orbitals, but the repulsive forces in two unbonded pairs results in 4 electron domains. The two unbonded pairs will take up even more room, exerting more of a repulsive force and altering the shape by creating a 104.5 degree angle and a bent molecule. This can be viewed as a tetrahedral shape with two of the bonds replaced by lone pairs. This can be represented by the symbol AX2E2, where A represents the central atom, X represents any atom bonded to the central atom, and E represents two unbonded pairs.

Bonding Theory

Compounds are formed by the forces of attraction between elements. These forces of attraction (bonding) occurs when an electron is transferred (ionic compounds) or shared (molecular compounds) between the atoms within in a compound.

Delocalized Electrons and Conductivity

Conductivity requires the presence of charged ions or particles to transmit the electrical current. Since negatively charged electrons are delocalized amongst positively charged cations, they pass the electric current across the metal from one atom to another, such that the electric current is transferred across the whole metal. For thermal conductivity, as metals are heated, the free valence electrons move more as they receive more kinetic energy. As they collide with adjacent particles, they transfer this energy, transferring heat in the process.

Delocalized Electrons and Properties of Metals

Delocalized electrons will also explain the conductive, malleable, ductile and lustrous nature of metals.

Polarity of Water and Hydrogen Bonding

Due to the strong polar nature of the water molecule, the hydrogen atoms within a water molecule can form hydrogen bonds with other oxygen atoms of other water molecules (coehsion), strongly electronegative atoms, or negatively charged anions, resulting in bonds that are stronger than regular dipole-dipole interactions. This property of the water molecule allows liquid water to dissolve many substances and act as a universal solvent.

Lone Pairs

Electron pairs that are not involved in the transfer or sharing of electrons, but are initially present as part of the atom, are referred to as lone pairs (unbonded pairs).

Trends in Electronegativity

Electronegativity increases from left to right (more protons) and from bottom to top (less distance).

Electrons and London Dispersion Forces

Electrons are continuously in rapid motion around atoms within compounds. As a result, temporary dipoles can be formed when electrons tend to briefly move towards one end of a molecule, creating a slightly negative end and a slightly positive end within an instant. These temporary dipoles can orient nearby molecules such that the slightly negative end repels electrons within an adjacent non-polar molecule and attracts the newly created slightly positive end (dipoles build on each other) Although these forces are very weak, when this happens over several molecules, the intermolecular force is significant enough to create attractions. Having more electrons increases the likelihood of temporary dipoles.

Energy Levels

Electrons exist within specific orbits or energy levels around the nucleus. Electrons in energy levels nearest the nucleus have the lowest energy, while energy levels farther away have more energy. Because of the positively charged nucleus, electrons in the first (lowest) energy level are more tightly held in the atom. Thus, it is only the outermost electrons (valence electrons) at the highest energy level that can be transferred and shared and are ultimately responsible for bonding.

Structural Digram

For a simple representation of what a compound may look like, a structural diagram may be drawn, whereby the bonding pairs are represented by dashes and lone electron pairs are not shown.

Determining Molecule Polarity

For complicated molecules, if the molecule is symmetrical (ie. all of the bonds within the molecule are identical) the molecule is most likely non-polar as the identical electronegativities cancel each other out. If all the bond polarities are identical to each other in opposite directions (same atom), they cancel each other out and result in a non-polar covalent bond. If the bond polarities are not identical (different atoms), additive vectors are possible, resulting in a polar molecule.

Molecular Compounds

Formed when two or more non-metals combine. Molecular compounds share electrons; they do not donate or borrow. There is an electrostatic attraction between two nucleus for one electron.

Hydrogen Bonding and Alcohols

Hydrogen Bonding is the reason for the high boiling points of alcohols. Alcohols contain a hydroxyl group (OH) and thus can have hydrogen bonding between the alcohol molecules, which require more heat and energy to break this attraction.

Permanent Dipoles and Hydrogen Bonding

Hydrogen bonding has the same principles of attraction as that of permanent dipoles. However, because of the existence of stronger attractive forces due to extreme dipoles (which occur as a result of high electronegativity for the lone electron of hydrogen, the "nakedness" (accessibility) of the hydrogen proton, and the lone (unbonded) pairs that increase the negativity of the dipole on N,O,F), these forces have been given a class of their own.

Hydrogen Bonding Stabalization

Hydrogen bonding is also responsible for the stabilization of other biological molecules. 1. For instance, hydrogen bonds also stabilize the double helix structure of DNA by holding the two DNA strands together. 2. Hydrogen bonds between hydrogen and nearby oxygen atoms in other amino acids can stabilize the secondary structure of proteins.

Polar Bonds and Polar Molecules

If a molecule contains a polar bond, is itself polar? For example, in a water molecule the horizontal vectors cancel out, but the vertical vectors are additive. Due to the additive nature of the vertical vectors, the water molecule is polar. However, in carbon dioxide, the horizontal vectors cancel out, even though the bonds themselves are polar, the molecule is non-polar covalent.

Valence Energy Level is Not Full

If the outermost energy level is not full, then the atom has a greater tendency to lose/gain electrons to make a full energy level (ie. is reactive).

Coordinate Covalent Bonds

In covalent bonds, each atom contributes/shares one electron to the electron pair to form the bond. However, coordinate covalent bonds (a.k.a dative covalent bond or dipolar bond) are formed when an atom contributes both of its electrons to form a bond rather than just one electron. Once formed, its strength and description are the same as the other covalent bonds within the compound. This can result in a compound that is charged as only the nucleus from the other atom or an additional electron may be transferred.

Resonance in General

In general, whenever electrons can be shared amongst three atoms, a resonance structure will exist. Because the electrons involved in bonding shift back and fourth amongst the three atoms, the strength of the bond is less than a double bond and is not considered a "fixed" double bond.

Crystal Lattice and Inter Forces

In the crystal arrangements, the cation (because of its positive charge) will be surrounded by the anions of other formula units, while the anion (because of its negative charge) will be surrounded by the cations of other formula units. This arrangement is very stable and and will remain regardless of whether the ionic compound is a big block or a fine powder. This organization occurs naturally and is also the reason for the solid state of ionic compounds (no room for ions to move into different states).

Inter Forces

Intermolecular and interionic forces are the attractive forces between molecules and formula units.

Bond Types and Electronegativity

Ionic Bonds: Metals have the lowest electronegativities, they do not attract valence electrons strongly and thus can lose them more readily to non-metals. As a result, metals transfer their valence electrons to non-metals to form an ionic bond. Molecular Compounds: Since elements that exist as diatomic molecules have some of the highest electronegativities, both atoms attract electrons strongly with equal strength such that there is no loss or gain of electrons. Different Non-Metal Atoms and Different E.N: The non-metal atom with higher electronegativity will attract electrons more strongly than the other atom, resulting in one end of the molecule being more negative and the other end being more positive (ie. a polar molecule). These bonds are known as polar covalent bonds.

Ionic Compound Structure and their Brittle Nature

Ionic compounds are held together through a transfer of electrons and this allows formula units to arrange themselves within a crystal lattice structure. If the structure becomes misaligned, then the like charges of the negative and positive ions will repel each other resulting in breaks along smooth planes.

Effect of Ionic Compound on Melting and Boiling Point

Ionic compounds consist of strong electrostatic attractions between positively charged cations and negatively charged non-metal anions. The electrostatic attraction is also responsible for arranging formula units of ionic compounds into highly ordered regular solid structures known as a crystal lattice. As a result, a lot of energy is required to break these electrostatic attractions between both formula units and compounds (intra and inter forces) to change the state, resulting in high metling and boiling points.

Ionic Compounds and Formula Units

Ionic compounds form when metals lose electrons (forming positively charged cations) to nonmetals which form negatively charged anions by gaining electrons. A single compound formed between metal and nonmetal ions is called a formula unit.

Lewis Formula

Is a type of diagram that consists of a Lewis dot diagram for a compound, and provides insight to the nature of bonding involved. Because of the octet rule, it is important to ensure all atoms, apart from the five exceptions, have a total of 8 valence electrons associated with them.

Metallic Bonding

Metals are solids and must have some attractive forces holding the metal atoms together (if not, they would be gases). They can transfer the few valence electrons they have to form ionic bonds but cannot accept enough electrons to fill their valences. To counteract this, all metal atoms share all valence electrons which are free to move from one atom to the next like a "sea", thus electrons are delocalized.

Ionic Bonding Process

Metals lose an electron and form positively charged ions called cations, while non-metals gain an electron and form negatively charged ions called anions. When a metal atom is reacted with a non-metal atom a transfer of electrons occurs such that the metal becomes a cation and the non-metal becomes an anion. The opposite charges of both attract to form an ionic compound.

Molecular Compounds and Molecules

Molecular compounds are formed when non-metals share valence electrons to fill the valence orbitals of both non-metals. A single compound formed between non-metals is called a molecule.

Molecular Compounds and Conductivity

Molecular compounds are neutrally charged as their electrons are involved in covalent bonding. As such, they are not free to conduct an electric current, even if the molecule is polar. Dissolution in water only disperses the individual molecules such that no electric charge can be conducted (no true charge).

Effect of Molecular Compound on Melting and Boiling Point

Molecular compounds consist of nonmetals that are covalently share electrons with each other (intra force) and interact with other molecules through Van der Waals forces (inter force). Covalent bonds within a molecule do not have to be broken (intra force) only the already weak van der Waals interactions have to be broken. As such, the melting and boiling points of molecular compounds are lower.

Hydrogen Bonding and Water Properties

One oxygen atom within a water molecule can bind up to 6 hydrogen atoms in liquid water at the same time. This is responsible for the fluid consistency of water and what gives snowflakes their characteristic six sided shape. This capacity for hydrogen bonding is also responsible for the high boiling point of water, allowing it to act as a heat sink (as more energy is required to break all of the hydrogen bonding.

Molecular Compounds and Solubility

Only intermolecular forces have to be overwhelmed to dissolve molecular compounds. Intramolecular forces do not need to be overwhelmed as molecular compounds disperse into molecules and not individual atoms.

Partial Charges and Polar Covalent Bonds

Partial charges within polar covalent bonds are represented with the symbol ₈ and its associated charge (₈⁻ or ₈⁺). Because these bonds contain a negative pole and a positive pole, they are referred to as bond dipoles.

States of Matter: Liquid and Gas

Particles then immediately form new bonds with other particles. This breaking and reforming of bonds result in particles continuously 'changing partners" resulting in the fluid state of a liquid. If the temperature is still increased, then the kinetic energy is also increased until there is enough energy to break all of the bonds holding the particles together. The particles now remain free of all other particles and are in the diffuse state of a gas.

Bonding Pairs

Since an orbital can have a maximum of 2 electrons, bonding electrons (valence electrons involved in bonding) are either shared or transferred to make bonding pairs of electrons to fill the orbitals.

Double or Triple Bonds

Since double and triple bonds occur within the same orbital plane as an existing "single bond", they function as one group while repelling electrons around a single atom. The forces of attraction are greater in double bonds since 2 electron pairs are involved in bonding, and more so in triple bonds since 3 electron pairs are involved in bonding. This means that the bond length (distance between elements involved in bonding) is shorter and much more rigid.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Since like charges repel each other, electron pairs within a compound will try to take a position as far apart as possible. Because lone pairs (unbonded pairs) are not engaged in bonds, they remain closer to the nucleus (attracted by only one nucleus in one direction), take up more room and exert greater repulsive forces than a bonding pair does, influencing the shape of the molecule The order of repulsion in greatest to least is : LP-LP>LP-BP>BP-BP

London Dispersion Forces

Since non-polar molecules do not have regularly positively and negatively charged dipoles, they do not interact in the same way as permanent dipoles or hydrogen bonding. All covalent compounds have London dispersion forces.

Lewis Symbols

Since only valence electrons are involved in bonding, Lewis proposed another method of representing bonding within compounds by drawing the element symbols, bonding electrons, and lone pairs. 1. Write the symbol of the element 2. Draw the valence electrons as a dot in each of the valence orbitals the element has, first singly and then adding dots to make electron pairs until all valence electrons are accounted for. 3. You can not have two lone pairs on opposite sides of each other.

Molecular Compounds and Forces

Since the nature of bonding within molecules in not due to a charged ion attraction, molecular compounds are not always arranged in the same form that ionic compounds are and can exist in other states as a solid, liquid, or gas. These forces are weaker than the charge attraction associated with ionic bonds, but the overall quantity of such forces can maintain molecular compounds in particular states and shapes. These forces consist of the following: 1. Dipole-Dipole Interactions 2. Hydrogen Bonding (naked Hydrogen) 3. London Dispersion Forces

Ionic Compounds and Conductivity

Solid ionic compounds do not conduct electricity because the compounds themselves are electrically neutral (cations and anions are held together as a compound, negating their charges) However, when in solution, interaction with polar water molecules can dissociate the compound into individual cations and anions, resulting in charged ions capable of conducting an electrical current. The solubility of ionic compounds may be reflective of the size of the ionic compound.

Shape of Sulfur Dioxide (bent - trigonal planar group)

Sulfur dioxide is formed through a coordinate covalent bond and is stabilized through resonance. Although bonding involves two orbitals within this compound, the unbonded pair of electrons will act "like a bond" by repulsing the other atoms, resulting in 3 electron domains. Just like trigonal planar molecules, the repulsive forces of 3 electron domains will push the peripheral atoms as far away from each other, but because there is an unbonded pair of electrons, they exert a greater repulsive force and take up more room, resulting in a 119.5 degrees angle and a bent molecule. This can be represented by the symbol AX2E, where A represents the central atoms, X represents any peripheral atom bonded to the central atom, and E represents a lone pair.

Discrepancy Between Melting and Boiling Point

The amount of kinetic energy required for all the bonds to be broken and undergo boiling to become a gas is significantly greater than the amount of energy required to melt a substance.

Orbitals

The area where electrons are likely to reside around the nucleus is referred to as an orbital, with each orbital having 0, 1, or 2 electrons. An energy level may have many orbitals depending on where it is located (s, p, d, f)

Electronegativity Trends

The attractive properties of an atom for valence electrons is referred to as electronegativity. The higher the EN value, the greater the atom's attraction for valence electrons. Fluorine has the highest EN (4.0), and EN increases as you go from metals (on the left) to non-metals (on the right). Noble gases have no electronegativity because they do not naturally form bonds.

Polarity and Shape

The bent and pyramidal shapes will always result in a polar molecule because the electronegativities will never all cancel out due to the lone pairs.

Bond Lengths and Resonance

The bonds in a benzene ring were found to be identical in length. This meant that each carbon was bound to each other by a bond and a half. This made the structure highly stable as a lot more energy is required to break the 1.5 bonds between any carbon (less distance = more attraction). This property of electrons being delocalized and shared amongst multiple atoms is called resonance.

Inter Forces and Ionic Compounds

The charges associated with the cations and anions within a formula unit will orient the formula units to organize into a highly ordered, regular arrangement - crystal lattice.

Delocalized Electrons and Lustre

The delocalized nature of the valence electrons allows for the concentration of these electrons at the surface, where they oscillate at a collective frequency. This oscillation reflects back light at a high enough intensity such that they appear lustrous.

Delocalized Valence Electron Attraction

The delocalized valence electrons act as a glue to maintain attraction of all the metal atoms together. Metals that contain more valence electrons can contribute more "glue" to the other metal atoms, resulting in a stronger metal (eg. group 1 vs. group 2 metals). Furthermore, metals that are smaller in size (ie. have a smaller atomic radius) will also form stronger metals as the distance between the valence electron and the positively charged nucleus is smaller, resulting in a stronger attaction.

Ionic Bond

The electrostatic attraction between a metal cation and a non-metal anion is called an ionic bond.

Factors Influencing the Strength of Temporary Dipoles

The greater the number of electrons, the higher the probability that temporary dipoles will be formed, and the stronger these dipoles will be since more electrons can create a stronger dipole. Since the number of electrons increases with the mass (due to an increase in the number of protons), the higher the mass, the stronger the London dispersion forces. As molecules increase in size (by gaining more energy levels), the lower the electronegativity and corresponding attraction of the valence electrons. As such, electrons can move more freely and are more likely to form temporary dipoles with greater frequency. Thus, the greater the size of the molecule, the stronger the London dispersion forces. Look at the states of the diatomic halogens: F2 (g), Cl2 (g), Br (l), I2 (s). The more linear the molecule, the greater the surface area for attraction, and hence the stronger the force between such molecules.

Alloys

The majority of metals are not pure but rather homogeneous mixtures (solutions) of different metals. These metals are called alloys and can have different properties than the individual metals themselves. Ex: 1. Steel is an alloy of iron and copper 2. Brass is an alloy of copper and zinc 3. Bronze is an alloy of copper and tin 4. Pewter is an alloy of tin, lead and copper.

Electronegativity Difference

The nature of a bond within a compound can be determined by calculating the electronegativity difference. The ∆E.N = Higher E.N - Lower E.N 1. If the ∆E.N is greater than 1.7, then the bond is ionic (assume transferred) 2. If the ∆E.N is less than 0.5, then the bond is slightly polar (similar E.N values) 3. If the ∆E.N is between 0.5 and 1.7 then the bond is polar covalent (partial charges). 4. If the ∆E.N is 0, then atoms are identicle and the bond is non-polar covalent (pure covalent)

Electronegativity (EN)

The nucleus contains protons (positively charged) that attract electrons (negatively charged). Different atoms have a different capacity for attracting electrons and forming bonds. The greater the number of protons in the nucleus, the greater the atom's attractive forces for electrons. The farther away the electrons are from the nucleus (higher energy-levels), the weaker their attraction to the protons in the nucleus is. The more electrons that exist between the valence electrons and the nucleus, the more the valence electrons are shielded from the attractive forces of the protons in the nucleus, through repulsion.

Effect of Charge on Melting and Boiling Points

The size of charge in ionic compounds and polarity in molecular compounds also affects the magnitude of the melting and boiling points. Higher charges increases the magnitude of the electrostatic forces while polarity increases dipole-dipole interactions, thereby increasing the melting and boiling point.

States of Matter: Solid

The state of a pure substance depends upon the strength of the attractive forces between the particles. In solids, the attractive forces are strong and the kinetic energy of the particles are not high enough to break the bonds that exist between them. If the temperature is continually increased, the kinetic energy is also increased until the energy is great enough to break the bonds between neighbouring particles.

Boiling Point

The temperature at which the state changes from a liquid to a gas is called the boiling point, and the amount of energy required for one mole of liquid to become gas is called the enthalpy of vaporization. The temperature will remain the same at the boiling point until the entire substance has vaporized. Temperature remains static during phase change.

Melting Point

The temperature at which the state changes from a solid to a liquid is called the melting point and the amount of energy required to melt one mole of a substance is called the enthalpy of fusion. The temperature will remain the same at the melting point until the entire sample has melted. Temperature is static during phase change.

Intra Forces

The types of attractive forces between the atoms within a compound (formula unit and molecule) are referred to as intramolecular or intraionic forces. In order to break the bonds between the formula units, the ions need to be dissociated.

Causes of Hydrogen Bonding

This bonding is due to a very positively charged dipole created when the hydrogen electron gets strongly attracted to the protons in another atom's nucleus, as well as the presence of lone (unbonded) pairs in that atom which push the hydrogen atoms away to create a highly negatively charged dipole. As such, only atoms with a high electronegativity and lone (unbonded) pairs are capable of hydrogen bonding with hydrogen. These atoms are O2, N2, and F2 (NOF)

Representing Polarity of a Molecule

This difference (and hence the polarity of a molecule) can be represented in the shape formulas by replacing "Y" for an "X" where: "Y" represents an atom bonded to the central atom that has a different magnitude and direction in the polarity of its bond with the central atom. "X" represents another atom bound to the central atom with a different polarity than "Y" "A" represents the central atom.

Resonance

Two structures are possible for Benzene (C6H6), and both structures will shift or resonate back and forth between each other constantly, lending stability to the structure. In fact, neither structure. actually exists but the molecule is actually a hybrid of both structures with three atoms sharing electrons. Because the electrons are not held or fixed to a particular atom (ie. carbon) they are considered delocalized.

Van Der Waals Forces

Van der Waals forces are the attractive or repulsive forces between molecules. These forces encompass dipole-dipole interactions and London dispersion forces.

Hydrogen Bonding

Water is a polar molecule with a bent (tetrahedral group) configuration. Hydrogen has only one electron in its valence orbital which is already engaged in a covalent bond with the highly electronegative oxygen atom in a water molecule, the rest of the hydrogen nucleus is mostly bare, creating a very positive end. The oxygen atom within a water molecule also contains lone pairs, which bend the hydrogen atoms away and make this end negatively charged.

Hydrogen Bonding and Freezing

When water freeze to form ice, each water molecule is bonded to four other water molecules. As a result, the water molecules are farther apart and the hydrogen bonds are longer, resulting in an arrangement that is more spaced out. In fact, in a water molecule, the oxygen atom is directly across its adjacent hydrogen atom, forming a straight line. This open arrangement causes ice to become less dense than liquid water, allowing it to float on the surface and act as an insulator in lakes during the winter.