Warmup: For each of the following, (1) determine if they will gain or lose electrons, (2) how many they will gain/lose, (3) write out the ionic symbol. Oxygen (#8) Potassium (#19) Calcium (#20)

New Objectives: Red, Yellow, Green Goal 1: I can identify all 3 types of chemical bonds. Goal 2: I know how to use subscritps and put together cations and anions to make an ionic compound. Goal 3: I know the difference between monatomic and polyatomic ions

3 types of bonds Ionic: very strong, creates solids. Polar: medium strength bond. Some solids, liquids, and gases Covalent: very weak bond. Some solids, liquids, gases

Bond type depends on how electrons are shared between atoms Electronegativity: An atom’s attraction for electrons High EN = strong attraction Low EN = low attraction

How to figure out bond type: You must calculate at the electronegativity difference between the two atoms Use the approximate guideline for electronegativity differences: 0.0 – < 0.5 = covalent (electrons shared nicely) 0.5 – < 2.0 = polar covalent (electrons being pulled from stronger atom) 2.0 – 4.0 = ionic (electrons stolen/transferred by stronger atom

Examples: H and Cl 2.1 – 3.0 Difference = 0.9 = polar C and H 2.5 – 2.1 Difference = 0.4 = covalent

Examples: Na and Cl 0.9 – 3.0 Difference = 2.1 = ionic Also notice that you have a metal bonded to a nonmetal! Follow this link for a video recap: https://www.youtube.com/watch?v=02Q352- Y7iU

Use these values provided to help: EN increases EN decreases

Atomic Size

Why? As you go across a period, the principal energy level (# of rings) remains the same. As the nuclear charge (# protons) increases, it draws those electrons in closer to the nucleus. Atomic size therefore decreases going across the periodic table, Atomic size increases when you move down the periodic table because more rings of electrons are being added.

Ionization Energy Energy required to remove an electron from a gaseous atom. This causes a +1 charge. The energy required to remove the first outermost electron is called the 1st ionization energy. The larger the atom is, the easier it is to rip away an electron. Electrons that are far away from the nucleus are easier to pull away.

Electronegativity Trends Increases Decreases

Ionic Bonding

Ionic Bonding Bonding of ions Ionic Bond Ionic Compound The electrostatic force that holds oppositely charged particles together in an ionic compound. One or more electrons are transferred from cation (metal) to anion (nonmetal or group of nonmetals). Ionic Compound Always formed between a metal and a nonmetal (or a group of nonmetals)

Compound Formation and Charge Want compounds to have a neutral charge Number of electrons lost = number electrons gained Example: CaF2 Calcium wants to lose 2 electrons to be like Argon. Fluorine can only gain 1 electron to be like Neon. This compound must use 1 Calcium and 2 Fluorine atoms.

Formula Unit The simplest ratio of ions represented in an ionic compound Overall charge of the formula unit is zero

Formula Writing Symbol of the cation is always written first Symbol of the anion is always written second Subscripts represent the number of ions present Example: What is the formula for an ionic compound containing magnesium and bromine?

Naming Ionic Compounds Recall: Cations keep the original element’s name Ex. A sodium ion is still called sodium Anions place “ide” at the end of their element’s name Ex. A chlorine ion is called chloride Since ionic compounds are made of cations (first) and anions (second) we will use the above rules to name ionic compounds Example: NaCl is called sodium chloride What would MgS be called? Magnesium sulfide

Polyatomic ions “poly” = many “atomic” = atoms Group of atoms bonded covalently that have an overall charge. Ex: PO4-3 = “phosphate” ion (5 atoms: 1 P atoms and 4 O atoms) Notice the ending of the name: -ate. This is NOT a monatomic anion from the periodic table You WILL NOT be required to know their formulas or charges. We have a chart for you.

You will be provided a LIST of polyatomic ions and their charges

Rules for using polyatomic ions If you need more than one polyatomic ion in a formula, you must group it together in parenthesis with a subscript outside. Ex: Try to make a balanced compound using magnesium and Chlorate. Remember: if the name of an ion ends in –ide, it came from the periodic table and is a monatomic ion. The end of this name is –ate If the ion ends in –ate or –ite…look it up on your list!!!

Some examples mixed up: Try these: Sodium oxide Calcium hydroxide (this is a polyatomic anion event though it ends in –ide) Strontium sulfite (anion ends in –ite…look it up on your list…it’s polyatomic) Barium nitrate Silver chloride Rubidium sulfate

Writing the Formula from the Name Find the elements that make up the compound Determine the charge of the ions Determine the ratio of the ions in the compound Write the formula, using () if you need more than one polyatomic ion. Example: What is the formula for calcium iodide?

Red? Yellow? Green? Goal 1: I can identify all 3 types of chemical bonds. Goal 2: I know how to use subscritps and put together cations and anions to make an ionic compound. Goal 3: I know the difference between monatomic and polyatomic ions