Chemical Bond An attraction between the nuclei and valence electrons of different atoms, which binds the atoms together When atoms form chemical bonds their valence electrons move around. This makes atoms more stable and determines the type of bond. Ionic Bond A chemical bond formed by the electrical attraction between positive ions and negative ions. When an atom steals an electron from an another atom those atoms then become bonded together because their strong opposite charge (positive and negative). Cation A positively charged ion An atom that had an electron stolen from it Cations are smaller than normal atoms because the positively charged nucleus has more power and pulls the electrons in closer.
Anion A negatively charged ion An atom that has stolen an electron from another atom Anions are bigger than normal atoms because the positively charged nucleus has less power and the electrons float further away. Covalent Bond A chemical bond formed by 2 or more atoms sharing their valence electrons
Ionic Bonds If the difference in electronegativity of two atoms is greater than 1.7 the bond is a ionic bond. Nonpolar-Covalent Bonds Type of covalent bond where bonded atoms share electrons equally. Bonds between atoms with an electronegativity difference of 0.3 or less are nonpolar-covalent bonds. Polar Covalent Bonds Type of covalent bond where bonded atoms share electrons unequally. The electrons are more strongly attracted to the more electronegative atom Bonds between atoms with electronegativity difference from greater than 0.3 to 1.7 are considered polar-covalent bonds.
Determining Bond Type 1. Find electronegativity values for bonded atoms and note which atom has the strongest value 2. Subtract the smaller value from the large value to find the difference 3. Determine the type of bond Chemical Compound Substance formed from 2 or more elements chemically bonded together in fixed proportions. Ex: Water is a chemical compound it is 2 elements (Hydrogen and Oxygen) and is always 2 Hydrogen atoms and 1 Oxygen atom Molecule Neutrally charged group of atoms that are held together by covalent bonds Can contain 2 or more atoms of the same element (oxygen gas) or two or more atoms of different elements (water) Molecular Compound Any chemical compound whose simplest units are molecules 2 or more molecules together
Molecular Formula Shows the number of each type of atom needed to form a single molecule Bond Energy The energy required to break a chemical bond and form neutral, isolated atoms Lewis Structure Diagram showing the bonding between atoms of a molecule and the unbonded (lone pairs) of valence electrons that may exist in the molecule
Single Bond A covalent bond in which one pair of valence electrons is shared between two atoms A single dash or two dots between two atomic symbols shows the single bond Double and Triple Bonds 2 or 3 pairs of valence electrons can be shared in a covalent bond 2 dashes or 4 dots between two atomic symbols show a double bond 3 dashes or six dots between two atomic symbols show a triple bond
Creating a Lewis Structure 1. Determine type and number of atoms in molecule 2. Write the electron dot diagram for each atom 3. Calculate the number of valence electrons available for bonding 4. Construct the Lewis Structure - if carbon is present put it in the middle, if not put the element with the lowest electronegativity in the middle (besides hydrogen). Then connect the atoms with electron pairs. 5. Add unshared pairs of electrons so each atom has a full set of valence electrons 6. Count the electrons in the structure to be sure there are the same number you calculated in step 3 Ionic Compound Compound composed of positive and negative ions that are combined so that the amount of positive and negative charge is equal Example: Salt = Sodium Chloride = Na + and Cl -
Formula Unit Simplest collection of atoms from which an ionic compound s formula can be established (Basically it's the ionic version of a molecule) Example: NaCl Example: CaF2 Monatomic Ion An ion consisting of a single atom Example: H +, K +, Mg +2 Example: O -2, S -2 Polyatomic Ions Ions made up of more than one atom Example: NH4+, NO2-, ClO-
Oxidation Numbers Oxidation number is equal to the charge an ion has after it has given or taken an electron. Example: When salt is made Chlorine (Cl) takes an electron from Sodium (Na). Chlorine becomes -1 and Na becomes +1. The -1 and +1 are oxidation numbers. Writing the Formula for Ionic Compounds Step 1: Identify the elements in the compound Step 2: Write out the oxidation number for the ions involved Step 3: Balance the oxidation number out between the elements Ex: Potassium and Oxygen Ex: Aluminum and Sulfur
Writing Formula with Polyatomic Ions Since polyatomic ions exist as a unit, when writing them in a formula never change their subscript. If more than one polyatomic ion is needed, place parentheses around the ion and write the subscript outside the parentheses. Oxyanion A polyatomic ion composed of an element bonded to one or more oxygen atoms More than one oxyanion exists for some non-metals, such as nitrogen and sulfur Ex: NO3 and SO4 Naming Oxyanions Beginning of the word is always the start of the non-oxygen atom. Atom with the greatest number of oxygen atoms has the ending -ate and the atom with the least number of oxygen atoms has the ending -ite. Naming Ionic Compounds 1. Name the cation followed by the anion. The cation always comes first in the formula. 2. For monatomic cations use the element name 3. For monatomic anions use the beginning of the element s name plus the ending -ide 4. If the cation has multiple oxidation numbers then show the correct oxidation number by using roman numerals. 5. If a compound contains a polyatomic ion then use the name of the polyatomic ion.
Naming Binary Molecular Compounds 1. The 1st element in the formula is named first using the entire element s name 2. The 2nd element in the formula is named using its beginning and replacing the ending with -ide 3. Prefixes are used to show how many atoms of each element are present