Chemistry- The Periodic Table and Bonding

family

a collection of elements from the same vertical group that have similar chemical properties. All of the members of a particular family have the same # of valence electrons

triple bond

a covalent bond in which two atoms share three pairs of electrons

Family of alkali metals (1A)

all of the atoms of elements in the extreme left-hand group have 1 valence electron. Include elements Li to Fr. Except for H, these elements are all extremely reactive. Physically, they are shiny, grayish-white metals. But, they melt more easily than the metals like iron or copper. They also tend to have lower densities than the more common metals.

Single bonds

covalent bonds in which one pair of electrons is shared b/w 2 atoms

Halogens (7A)

group of elements alongside the noble gases. The halogens are very reactive. The elements are quite physically distinct from one another. F and Cl are greenish-yellow, toxic gases; Br is a brown liquid at room temp; and Iodine (I) is a grayish-purple solid. All have 7 valence electrons, so they have similar chemical properties. Elements F to At

Main types of bonds

ionic, covalent (nonpolar covalent and polar covalent), and metallic

Electrostatic force

the attraction b/w a positive charge and a negative charge; this force is very strong. The strength of an ionic bond gives ionic compounds their high melting points, hardness, and other physical properties.

Binary ionic compound (contains only 2 elements) naming

the cation in the compound keeps its name, and the anion changes its ending to -ide. When a polyatomic ion is present as part of the compound, it keeps its name.

Nomenclature

the study of naming chemical compounds. It's based on the type of bonding

Lone pairs

unbonded electron pairs

Covalent vs Ionic chemical reactions

when a covalent substance undergoes a phase change, bonds are NOT breaking as they would in an ionic substance

Summary of Molecular Shapes

-# of Electron Pair Sites Around Central Atoms; # of Lone Pairs Around Central Atom; Shape; Bond Angles w/ Central Atom (assuming atoms of the same element surround the central atom); Example -4, 0, Tetrahedral, 109.5, CCl₄ -4, 1, Trigonal Pyramidal, approximately 107, NH₃ -4, 2, Bent, approximately 107, H₂O -3, 0, Trigonal Planar, 120, NO₃⁻ -3, 1, Bent, approximately 116, SO₂ -2, 0, Linear, 180, CO₂

Naming Transition Metals

-B/c transition metals can form cations w/ multiple charges, the charge of the cation must be specified in the name of the compound using Roman numerals. -Ex) copper (I) sulfate is Cu₂SO₄ copper (II) sulfate is CuSO₄

Elements that violate the octet rule

-Beryllium (Be) atoms are stable w/ 4 valence electrons. When Be is the central atom the molecule is linear. -Boron (B) atoms strive to gain 6 valence electrons. When B acts as the central atom in a molecule, the shape is generally trigonal planar.

Coulomb's Law

-E = (k * q₁ * q₂)/r, where E is the amount of energy (needed to remove an electron from an atom), k is a constant, q₁ and q₂ are the charges on each ion, and r is the length of the bond. -Allows us to determine how much energy is present in any ionic bond -A shorter bond length leads to a greater bond/lattice energy -Greater charge and shorter distance lead to greater energy

Ionic Formulas

-Oxygen group atoms typically have 2- charge -Halogen group typically have -1 charge -Nitrogen group typically have 3- charge -Aluminum group have +3 charge -In group 4, C and Si like to form covalent bonds but Pb and Sn typically form ions w/ a +4 charge -Alkaline earth metals have +2 charge -Alkali metals have +1 charge -Transition metals have various charges. Some lose electrons to form cations but the # of electrons they lose can vary and they often form ions with multiple possible charges. -To determine formula of an ionic compound, the total charge on an ionic compound must be 0. The cation always comes first in the formula of Na₂S.

Bond Energies

-a short bond length will lead to greater energy -the more bonds there are b/w 2 atoms, the shorter the bond length is going to be. So a triple bond is the shortest and strongest type of covalent bond while single bonds are the weakest. -If the # of bonds is identical, we can look at the size of the atoms involved in the bond to determine the bond length. A smaller atom leads to a smaller bond and therefore greater bond energy/strength.

Nonpolar vs Polar Molecules

-any diatomic molecule that has a polar bond is polar. Ex) CO. -It's possible for a molecule to contain polar bonds, but itself, be nonpolar (the individual bond polarities cancel each other out.). Ex) Methane (CH₄) -Molecules that consist of 3 or more atoms are generally polar unless... -... if the central atom has no lone pairs and is surrounded by atoms of one element, then the molecule will be nonpolar Ex) CO₂. -any diatomic molecule that has a nonpolar bond is nonpolar. Ex) all elemental diatomic molecules like Cl₂, N₂, and O₂.

Molecular Shapes

-assume the first atom in the formula is the central atom of the structure (unless its H, which is never a central atom) -using dots to indicate the valence electrons of each atom, surround the central atom w/ the others, trying to give each atom an octet. Remember hydrogen needs only 2, not 8, valence electrons to be satisfied. It's important to realize that electrons shared b/w 2 atoms count toward the total for both. -to determine the shape of the molecule, you must consider the # of sites in which valence electron pairs surround the central atom. The # of total electron pair sites and # of lone pairs will dictate the molecule's shape.

Nonmetals

-elements that tend to gain or share electrons when they bond -are usually poor conductors of heat and electricity -some such as Sulfur and Phosphorus are solids at room temp -are dull, brittle, and melt easily aka low melting point (diamond -non metal carbon- is an exception) -a few nonmetals like oxygen and fluorine are gases -nonmetal bromine is a liquid at room temp -physical characteristics of nonmetals vary -groups 4-8

Semimetals (metalloids)

-have some of the physical characteristics of both metals and nonmetals -For ex, Si is shiny like a metal but brittle like a nonmetal -lie b/w metals and nonmetals on the periodic table -can either gain, lose, or share electrons in a bond -B, Si, Ge, As, Sb, Te, Po

Reactivity

-how easily an atom gains or loses electrons -increases up and right (more likely to gain electrons) and down and left (more likely to lose electrons, except for noble gases) -elements in the middle of the table are mostly unreactive

Metallic Character

-how easily an atom gives up an electron in a bond to form a positive ion -as you move from left to right across a period, it decreases -as you move from top to bottom down a group (toward francium) it increases -decreases across the table; increases down the table

Polyatomic ions

-ions that are created of multiple elements that stay bonded together and act as a single unit when forming ionic compounds. -if there are multiple polyatomic ions present in a compound, parentheses must be used to show that. *- NO₃⁻ Nitrate - SO₄²⁻ Sulfate - CO₃²⁻ Carbonate -OH⁻ Hydroxide - PO₄³⁻ Phosphate - NH₄⁺ Ammonium*

Electronegativity

-the amount of "pull" that an atom's nucleus exerts on another atom's electrons when it's involved in a bond. -As you move across the periodic table from left to right, electronegativity increases. As you move down a group, the electronegativity decreases. Electronegativity increases as you move towards F -Same trend as ionization energy

Ionization energy

-the amount of energy required to remove an electron from an atom -ionization energy increases as you move from left to right across the periodic table; gets harder to remove an electron from the atom -As you move from top to bottom through a group, ionization energy decreases. -ionization energy increases from left to right across the periodic table as you move up through a group (it increases as you move towards F) -As you remove additional electrons past the first, the atom (now ion) will get progressively smaller. So, removing each successive electron will become more difficult and the remaining electrons are closer to the nucleus and more attracted to it. -The second ionization energy from an atom is greater than the first, the third would be greater than the second, etc. -get BIG jumps in ionization energy when you start removing electrons from a lower principle energy level.

Atomic radius

-the distance from an atom's center to the edge. The larger the atom, the greater its radius -As you move across the periodic table from left to right, atomic radius decreases. -As you move down a group (towards francium), atomic radius increases.

Lattice (binding) energy

-the lattice energy of an ionic solid is a measure of the energy required to completely separate a mole of a solid ionic compound into its separate ions. It's the bond energy present in an ionic compound -The higher the lattice energy, the stronger the ionic bond. -The greater the lattice energy present in an ionic substance, the higher the melting point of that substance will be.

Metals

-usually shiny and are good conductors of heat and electricity -many are malleable -often ductile (can form wires, for ex Cu) -all metals are solid at room temp, except for Mercury (a liquid) -tend to give up electrons when they bond -Active metals (group 1 and 2): the reactive metals of the s area. Active metals also tend to have lower melting points than transition metals. -Transition metals (b group and period 7): the rest. they are quite diff from active metals. They are generally harder, more difficult to melt, and less reactive than active metals. They include elements in the d and f areas. Include elements like iron, copper, gold, and silver. Many compounds that contain a transition metal are intensely colored.

Metallic Bond

-when 2 metals bond -the metal atoms donate valence electrons to become cations. These valence electrons are not directly transferred to another atom, but instead they move about freely throughout the sample, producing an attractive force that keeps the metal cations in place. -the behavior of these free electrons is referred to as a "sea of electrons" -b/c of the motions of free electrons, metals are typically good conductors of electricity and heat. -Ex) Cu, Ag, Fe

Covalent bond

-when 2 nonmetals bond -2 atoms share electrons to achieve a stable octet

Nonpolar covalent bonding

-when 2 nonmetals share electrons equally -can only occur b/w atoms with identical electronegativity values, so the only truly nonpolar covalent bonds are those present in molecules made up of one type of atom (like O₂)

Ionic Bond

-when an atom in a bond gives up 1 or more electrons to the atom it bonds with. -generally form b/w atoms that differ greatly in their electronegativity values. -the atom that gives up an electron becomes a positively charged ion and the one that accepts the electron becomes negatively charged. The + charged ion attracts the - charged atom and this draws the 2 atoms together and results in the release of energy. *when a metal and nonmetal bond, the result is an ionic bond, in which the atoms are held together by an electrostatic attraction b/w a positive and a negative ion* -substances that are held together by ionic bonds are solids at room temp and atmospheric pressure. Ionic solids are characterized by their hardness, brittleness, and high melting points. Ionic solids can't conduct electricity b/c their ions have very restricted movement. But, if an ionic solid is melted, its ions are freer to move and the substance can conduct electricity. -Ex) NaCl. Na gives its outermost electron to become Na⁺ and Cl receives it to become Cl⁻. The Na⁺ ion and Cl⁻ ion now have 8 electrons in their outermost shells

polar covalent bond

-when in a covalent bond, certain atoms have a partial positive charge and others have a partial negative charge. -one atom hogs the electrons b/c it has a greater electronegativity than the other -Ex) In H₂O, the hydrogen atoms have a partial positive charge and the oxygen has a partial negative charge b/c oxygen has a greater electronegativity

Covalent nomenclature

-when naming binary covalent compounds use these prefixes: one = mono two = di three = tri four = tetra five = penta six = hexa seven = hepta eight = octa -when naming the compound, how many of each atom present is represented by the appropriate prefix. Only exception is when the first element has only one atom present, you don't use the mono prefix

Lewis dot diagram

1. count the # of valence electrons in each atom and add them up 2. draw a skeletal structure of the molecule with the least electronegative atom in the center 3. create a single bond (shared electron pair) connecting the central atom to each terminal atom 4. Add lone pairs around each terminal and central atom until each atom has 8 total electrons (except hydrogen. which only needs 2) 5. count up the total # of electrons in the structure. If they equal the total # of valence electrons available (calculated in step 1) your structure is correct. If you have more assigned electrons than valence electrons, you need to shift some lone pairs over and create double or triple bonds.

double bond

A covalent bond in which two pairs of electrons are shared between two atoms

Active metals

B/c of their highly reactive nature, elements of the alkali and alkaline earth families are collectively known as active metals.

Noble gas family (8A)

Elements He to Rn. All elements in the noble gas family (including He) are very unreactive (inert). They all have 8 valence electrons except for He.

The family of alkaline earth metals (2A)

Group to the right of alkali metals from Be to Ra. They have 2 valence electrons. They are less reactive than the alkali metals but more reactive than common metals like iron and copper. They look very similar to alkali metals.