Excelsior Chemistry Unit 3 (Test A Only)

STUDY
PLAY
Valence Electrons
Electrons in the highest occupied energy level of an element's atoms. These are usually the only electrons used in chemical bonds. You can determine the number of valence electrons in a representative element by looking at the group number. So an element from Group 4 like Carbon means that it has 4 valence electrons.
Electron Dot Structure
Diagrams that show the valence electrons as dots. The electron dot structures show that elements of the same group have the same number of electron dots and valence electrons.
The Octet Rule
An octet means 8. Noble gases are stable and they all have 8 valence electrons. The octet rule means that elements try to form so that there are 8 valence electrons (like noble gases). In order to follow this rule, atoms of metals tend to lose their valence electrons, leaving a complete octet in the next -lowest energy level. Atoms of non-metals tend to gain electrons or share electrons with other non-metals to achieve a complete set. So Carbon, having 4 valence electrons already, would need to add 4 more in order to make the 8 needed to be stable like a noble gas.
Metals and the Octet Rule
Metal atoms tend to lose their valence electrons leaving the next level down as a full octet.
Non-metals and the Octet Rule
Non-metal atoms tend to gain valence electrons, or share them in order to complete a full octet.
Cations
Cations are created by atoms losing valence electrons (which are negatively charged) leaving only protons (which are positively charged) and thus creating a positively charged ion called a cation. For example, the element Sodium has 1 valence electrons and it loses an electron in order to achieve the electron configuration of a noble gas. It's new configuration after losing the electron is level 1s is full with two electrons, level 2s is full with 2 electrons and level 2p is full with six electrons. Therefore, 1s^2 2s^2 2p^6
Anions
Anions are created by atoms gaining valence electrons (which are negatively charged) leaving more electrons than protons (which are positively charged) and thus creating a negatively charged ion called a anion. For example, the element Fluorine has 7 valence electrons and it adds an electron in order to achieve the electron configuration of a noble gas. It's new configuration after gaining the electron is level 1s is full with two electrons, level 2s is full with 2 electrons and level 2p is full with six electrons. Therefore, 1s^2 2s^2 2p^6
Halide ions
The ions produced when halogens (including chlorine) gain electrons. All halide ions have a charge of -1 because they have 7 valence electrons and need only one electron in order to become stable like a noble gas.
Ionic compounds
Compounds composed of cations or anions. They join together to form neutral compounds. The total positive charge of the cations equals the total negative charge of the anions. Most ionic compounds are crystalline solids at room temperature and generally have high melting points.
Ionic bonds
The electrostatic forces that hold the ions in ionic compounds together. Salt (Sodium Chloride) provides a simple example of a compound held together by an ionic bond.
Chemical Formula
Shows the composition (kinds and numbers of atoms) for the smallest unit of a substance. NaCl is the chemical formula for Sodium Chloride or table salt.
Formula Unit
The lowest whole number ratio of ions in an ionic compound. Formula units take into account that the overall charge must be neutral. In table salt, or Sodium Chloride, these combine in a 1:1 relationship so their Chemical Formula is NaCl (one Sodium with one Chlorine). For Aluminum Bromide, the Aluminum has a charge of 3+ and the Bromine has a charge of 1-. Therefore, in order to keep the balance at 0, you would need three Bromines with one Aluminum or AlBr3.
Properties of Ionic Compounds
Most ionic compounds are are crystalline solids at room temperature with high melting points and can conduct an electric current when melted or dissolved in water.
Coordination Number
The number of ions of opposite charge that surround an ion in a crystal.
Metallic Properties
Metals are made of cations. The valence electrons of metal atoms can be modeled as a sea of electrons. They can move and drift freely throughout the metal. Metals are good conductors as a result. Metals are also ductile and malleable.
Metallic bonds
Consist of the attraction of the free-floating valence electrons to the positively charged metal ions.
Crystalline Structure of Metals
Metal atoms are arranged in ordered and compact patterns. They can be body-centric (each atom has 8 neighbors), face-centric (12 neighbors), or hexagonal (also 12 neighbors). These neighbors also serve as the coordination number for each type of pattern.
Alloys
Mixtures composed of two or more elements at least one of which is a metal. They are often stronger and have better properties than the single metal alone. The most important alloy is Steel.
Covalent Bonding
The atoms held together by sharing electrons are joined by a covalent bond.
Molecule
A neutral group of atoms joined together by covalent bonds. Molecules consisting of two atoms are called diatomic molecules.
Diatomic Molecule
A molecule consisting of two atoms. All halogens form diatomic molecules. The same happens with
Molecular compound
Just like it says, this is a compound composed of molecules. These compounds tend to have relatively lower melting and boiling points than ionic compounds.
Molecular Formula
The chemical formula of a molecular compound. It shows how many atoms of each element a molecule contains.
Octet Rule in Covalent Bonding
In covalent bonds, electron sharing usually occurs so that atoms attain the electron configuration of 8 valence electrons in the outermost level, which is similar to noble gases.
Single Covalent Bond
The bond between two single atoms held together by sharing a pair of electrons. This is represented by a single dash or two dots.
Structural Formula
Represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms. Ammonia is a molecule of gas with one unshared pair of electrons. Ammonia is comprised of one Nitrogen and three Hydrogens. It's structural formula has a Nitrogen with either side of the Hydrogen and one at 90 degrees from the others.
Double or Triple Covalent Bonds
Atoms form these bonds if they share two (Double) or three (Triple) pairs of electrons in order to fill their outer most level with valence electrons. Oxygen tends to form double bonds, and Nitrogen forms triple bonds.
Coordinate Covalent Bonds
A covalent bond in which one atom contributes both bonding electrons in a single covalent bond. In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms. Hydronium ions form by the addition of a hydrogen ion to a water molecule. Oxygen must share a pair of electrons with the added hydrogen ion to form a coordinate covalent bond.
Polyatomic ions
A tightly bound group of atoms that has a positive or negative charge and behaves as a unit. They usually contain both covalent and coordinate covalent bonds as well as ionic bonds.
Bond Dissociation Energy
The energy required to break a covalent bond. The higher the energy required, the stronger the bond. Triple bonds have a dissociation energy of 908 kJ/mol, double bonds have 657 kJ/mol, and single bonds have 347 kJ/mol. That means that triple bonds are the strongest bonds, and that single bonds are the weakest bonds.
Molecular Orbitals
When two atoms combine to form a molecule, their atomic orbitals overlap to produce these which apply to the entire molecule instead of just the atom itself. Each atomic orbital is filled if it contains two electrons, and the same is true of molecular orbitals.
Bonding Orbital
The molecular orbital that can be occupied by two electrons of a covalent bond.
Sigma Bond
When two atoms bond with covalent bonding, their atomic orbitals combine to create molecular orbitals. If the new molecular orbital goes around both the nuclei in a symmetric pattern, a sigma bond has occurred.
Pi Bonds
When two atoms bond with covalent bonding, their atomic orbitals combine to create molecular orbitals. If the atomic orbitals were side by side, the new molecular orbital will form above and below the bond axis and a pi bond has occurred.
VSEPR Theory
According to this theory, the repulsion between electron pairs cause molecular shapes to adjust so that the valence electron pairs stay as far away from each other as possible creating a 3D molecule. In a water molecule, the two bonding pairs and the two unshared pairs of electrons form a tetrahedral arrangement around a central oxygen making the water molecule flat but bent at a 105 degree angle.
Hybrid Orbitals
This provides information about both molecular bonding and molecular shape. Using hybridization, atomic orbitals mix to form the same total number of equivalent hybrid orbitals. Single bonds hybrid orbitals form a tetrahedral angle of 109.5 degrees. Double bonds hybrid orbitals form a line with two angles on the ends at 120 degrees. Triple bonds hybrid orbitals form a linear molecule.
Van der Waals Forces
The collective name of the the two weakest attractions between molecules, dipole interactions and dispersion forces.
Dipole interactions
Occur when polar molecules are attracted to each other.
Dispersion Forces
The weakest of all, and caused by the movement of electrons as they move throughout the molecule.
Hydrogen Bonds
The attractive forces in which hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared pair of another electronegative atom. The attraction between the hydrogen of one water molecule and the oxygen of another water molecule forms a relatively strong attraction called a hydrogen bond.