Electron Configuration and Periodic Properties Atomic Radii The size of an atom is defined by the edge of its orbital Since this boundary is fuzzy, the radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together

Atoms tend to get smaller as you move across a period due to the increased positive charge They get larger as you move down a group due to the increasing energy levels occupied

Ionization Energy Ionization energy is the energy required to remove one electron from a neutral atom Made on isolated atoms in the gas phase In general, ionization energies of the main group elements (s&p) increase across a period Generally decrease down a group

With sufficient energy, electrons can be removed from positive ions as well as from neutral atoms The energies are referred to as the second ionization energy, third ionization energy, and so on These energies generally increase due to the stronger effective nuclear charge There are large jumps in energies when stable arrangements are ionized (in particular- the noble gas configurations)

Electron Affinity The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity Atoms that release energy have a negative affinity (they want the electron) Atoms that require energy to “force” the electron on them have a positive affinity (they will lose the electron spontaneously)

The halogens gain electrons most readily The p group elements generally become more negative as you move across a period (again exceptions caused by stable electron arrangements) The trends in groups are not as regular (competing increased nuclear charge and atomic radius) Generally the size predominates

For an isolated ion in the gas phase, it is always more difficult to add a second electron to an already negatively charged ion Second affinities are therefore always positive Ions like Cl -2 never occur

Ionic Radii A positive ion is known as a cation Caused by the loss of electrons The remaining electrons are drawn closer to the nucleus by the unbalanced charge A negative ion is known as an anion Formed from the addition of extra electrons The electrons are not drawn as tightly as they were before the addition

The metals on the left tend to form cations, while the nonmetals on the upper right tend to form anions Cationic radii decrease across a period due to increasing nuclear charge Anionic radii (starting w/ group 15) decrease across a period Ionic radii tend to increase down a group

Valence Electrons Chemical compounds form because electrons are lost, gained, or shared between the outermost energy levels of atoms (the inner electrons are too tightly held These available electrons are called the valence electrons For the main group elements these are in the s & p shells

Electronegativity Valence electrons hold atoms together In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons Fluorine is assigned a number of 4.0

Electronegativities tend to increase across each period Electronegativities tend to either decrease down a group or remain about the same Noble gases do not form many compounds and may not have values

Properties of the d and f block elements The properties of the d block elements vary less and with less regularity than those of the main group elements Both the outer s and the d electrons are available to interact with their surroundings The atomic radii of the d block elements generally decrease across a period The d electrons shield the outer electrons The electrons repel each other

The f block elements behave in a similar way Ionization energies generally increase across a period for d & f block elements In contrast, they generally increase down a group because the electrons available for ionization in the outer s level are less shielded (incomplete d shell) from the increasing nuclear charge

Ion formation in the d & f block elements follows the reverse order of electron configuration For d block, although electrons are being added to the d, they are removed from the outer s first (most d block elements therefore form +2 ions) The d & f block elements all have similar electronegativities Follow general trend

Chemical Bonding

Introduction to Chemical Bonding A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds atoms together Atoms bond because it decreases their potential energy, creating more stable arrangements of matter

Chemical bonding that results from the electrical attraction between large numbers of cations and anions is called ionic bonding Covalent bonding results from the sharing of electrons pairs between two atoms In a purely covalent bond, the shared electrons are “owned” equally by the two bonded atoms

Bonding is rarely purely ionic or covalent Electronegativity is a measure of an atom’s ability to attract electrons The degree of ionic or covalent character is determined by calculating the difference in electronegativity

The  indicates a partial charge

Covalent Bonding and Molecular Compounds A molecule is a neutral group of atoms that are held together by covalent bonds Individual unit capable of existing on its own May consist of two or more atoms A chemical compound whose simplest units are molecules is called a molecular compound

A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts A diatomic molecule is a molecule containing only two atoms

A balance is reached between the attractive forces and the repulsive forces between the nuclei and electrons. This results in the most energetically stable arrangement.

In a covalent bond, the electrons orbitals can be pictured as overlapping (the electrons are free to move in either orbital) The distance between two bonded atoms at their minimum potential energy is the bond length The atoms will vibrate a bit

The difference between the potential energy zero level (separate atoms) and the bottom of the valley (bonded atoms) is the bond energy that is released when the bond is formed It is also the energy required to break a chemical bond and form neutral isolated atoms Atoms tend to acquire noble gas configurations when bonding

Octet rule: Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level

There are exceptions to the octet rule Boron: In BF 3, boron will share its three valence electrons and acquire a total of 6 When some elements combine with the very electronegative atoms of F, O, and Cl, an expanded valence that involves electrons in the d orbitals occurs

Electron-dot notation is an electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol

Electron dot notations can also be used to represent molecules A shared pair of electrons is drawn between two atoms, an unshared pair is a pair of valence electrons that belongs exclusively to one atom and is not involved in bonding H:H

A shared pair of electrons is often represented with a dash The are called Lewis structures A structural formula indicates the kind, number, arrangement, and bonds, but not the unshared pairs of atoms in a molecule

A single bond is a covalent bond produced by the sharing of one pair of electrons between two atoms A double covalent bond is produced by the sharing of two pairs of electrons between two atoms A triple covalent bond is a bond produced by the sharing of three pairs of electrons between two atoms Double and triple bonds are referred to as multiple bonds

C, N, and O can have multiple bonds H can have only one bond

Resonance structures cannot be correctly represented by a single Lewis structure Ozone Once thought to split time between two structures Experiments show that bonds are equivalent ( average of two bonds)

Not all covalent compounds are molecular Some are continuous 3 dimensional networks of covalently bonded atoms Called covalent-network bonding

Ionic Bonding and Ionic Compounds An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal Most are crystalline solids The formula simply represents the simplest ratio of ions that give neutrality of charge – called a formula unit

To compare bond strengths in ionic compounds, chemists compare lattice energies Lattice energy is the energy released when one mole of an ionic crystalline solid compound is formed from gaseous ions

The attraction between positive and negative ions is generally very strong In molecular compounds, the covalent bonds are also very strong, but the intermolecular attractions are generally much weaker than ionic attractions The melting point, boiling point, and hardness of a compound depend on how strongly these basic units are attracted to each other

Many molecular compounds melt at low temperatures, while many ionic compounds have high MP and BP Ionic compounds are brittle (a shift in layers can cause a strong repulsive force)

As a solid, ions cannot move, so ionic compounds are not conductors In the molten or aqueous state, they are free to move and are conductors A charged group of covalently bonded atoms is known as a polyatomic ion

Metallic Bonding Chemical bonding is different in metals than in ionic, molecular, or covalent- network compounds The valence electrons are highly mobile For most metals, the highest p orbitals are vacant (and often some d orbitals as well) In metals, these vacant orbitals overlap The electrons then can roam freely throughout the metal (delocalized)

The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons is called metallic bonding As a result, metals have high electrical and thermal conductivity Since they have many orbitals separated by extremely small energy differences, they can absorb a wide range of frequencies (and radiate them back) – causes shiny appearance

Metals are malleable (hammered or beaten into shapes) and ductile (drawn into a thin wire) Caused by uniformity of bonding throughout the metal Metallic bond strength varies with the nuclear charge and the number of electrons in the metals electron sea Reflected in the heat of vaporization

Molecular Geometry Molecular properties depend not only on the bonding of atoms but also on molecular geometry The polarity of each bond, along with the geometry of the molecule, determines the molecular polarity There are two theories to explain geometry VSEPR hybridization

VSEPR Theory Stands for valence-shell, electron-pair repulsion States that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible Diatomic atoms are linear

The number of bonds determines the bond shapes

If the central atom has both shared and unshared electrons, the unshared electrons must be accounted for also They also take up space around the atom

Double and triple bonds are treated in the same way as single bonds Polyatomic ions are treated in the same way as molecules

Hybridization VSEPR is useful for explaining the shapes of molecules- but it does not reveal the relationship between a molecule’s geometry and the orbitals occupied by its bonding electrons Hybridization is the mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies

Methane (CH 4 ) is a tetrahedron How does carbon (outer shell 2s 2 2p 2 ) form four equivalent bonds? The 2s and three 2p orbitals hybridize to form 4 equivalent hybrid orbitals called sp 3 The orbitals all have energy that is greater than the 2s but less than the 2p

Hybrid orbitals are orbitals of equal energy produced by the combination of two or more orbitals on the same atom Explains many Group 15 & 16 elements

The linear geometry of BeF 2 can be explained by the hybridization of one s and one p orbital (called sp hybrid) BF 3 is trigonal planar Involves one s and two p orbitals Called sp 2 hybrid

Intermolecular Forces As a liquid is heated, the kinetic energy of its particles increases At the boiling point the energy is sufficient to overcome the forces of attraction between the liquid’s particles Boiling point is a good measure of the force of attraction between particles of a liquid (higher = stronger) The forces of attraction between molecules are known as intermolecular forces

The strongest intermolecular forces exist between polar molecules Polar molecules 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

Dipole The dipole direction is from the positive to the negative side Represented by an arrow with the head pointed toward the negative pole and a crossed tail at the positive pole

The negative region in one polar molecule attracts the positive region in adjacent molecules The forces of attraction between polar molecules are known as dipole-dipole forces Short range forces only

Shows in the difference between the boiling point of BrF (-20 o C) and that of F 2 (-188 o C)

For molecules containing more than two atoms, molecular polarity depends on both the polarity and the orientation of each bond

A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons This results in a short range intermolecular force that is somewhat weaker than the dipole-dipole force This accounts for the solubility of nonpolar O 2 in water

Some hydrogen containing compounds have unusually high boiling points This is explained by the presence of a particularly strong type of dipole-dipole force In compounds containing bonds between hydrogen and fluorine, oxygen, or nitrogen the large electronegativity difference makes them highly polar The small size of the hydrogen allows it to come very close to the unshared pair of electrons on 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 is known as hydrogen bonding

Hydrogen bonds are usually represented by dotted lines H 2 S boils at -61 o C while water boils at 100 o C

Since electrons are constantly moving, temporary dipoles can be created The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles are called London dispersion forces Fritz London proposed in 1930

London dispersion forces Act between all atoms and molecules They are the only intermolecular forces acting among noble gas atoms and nonpolar molecules (low boiling points) Since they are dependent on electron motion, they increase with the number of electrons (increase with increasing molar mass)