Groups of Elements Alkali Metals Alkaline Earth Metals Transition Metals Halogens Noble Gases
Trends in the periodic table Groups (families) They all have the same # of outer electrons= VALENCE ELECTRONS Periods They have valence electrons in the same highest energy level
Periodic trends Periodic = happens according to a repeating pattern Periodic table has repeating pattern of valence e- configurations Leads to periodic trends: Atomic Radius Ion Formation and Ion size Ionization Energy Electronegativity
Periodic trend: Atomic Radius Size of atoms depends on: Number of energy levels Number of protons Down a column: Across a row:
Atomic radius increases down a group. The number of energy levels increases down a group Each subsequent energy level is further from the nucleus. n=1 n=2 n=3 n=4 n=5
Atomic radius decreases across a row Across a row valence electrons are in the same energy level (n=3) Positive charge of nucleus is partially cancelled out by negative charge of core electrons =electron shielding With each additional proton, there is a stronger force pulling the electrons closer to the nucleus. This results in a smaller atomic radius. Electron shielding in Na vs. Cl
Ionization Energy Can you explain this trend? Click HERE for a tutorial.HERE
Ionic Radii What happens to the size of the atomic radius when an electron is removed? Is an anion bigger or smaller than its neutral atom? Why?
Radii of neutral and charged atoms in pm (10 -12 m) Purple models represent neutral atoms; Red=cations; Blue=anions
Check your understanding Which atom would be larger, and why? N or O K or Rb Ne or Na Which atom would be more difficult to turn into a cation (higher IE), and why? Li or Be Mg or Ca F or Ne
Electronegativity Electronegativity – is the ability of an atom to ATTRACT electrons when the atom is in a compound. (pg 363) Credit to Linus Pauling
Ionization & Ionic Compounds Unit 3 Chemistry 1 Spring 2012
What’s an ion? Ion = Cation Anion Atoms form ions when they are able to give away or accept electrons. The most likely type of ion an atom will form is based on the number of valence electrons it starts with.
Comparing cations and anions Cations Positive charge Valence e- removed Size vs. atom: Na Na + Cu Cu 2+ Anions Negative charge Valence e- added Size vs. atom: O O - O 2-
Ionization-how does it happen? How do ions form? Will it be endothermic or exothermic? What happens next?
Specific ions form for each element due to its electron configuration Elements with electrons that completely fill the s and p orbitals of the highest energy level do not react. Elements without “full” outer energy levels take, give up, or share electrons (chemical bonding). Vocabulary: OCTET RULE ISOELECTRONIC
Cation formation How many valence electrons does Na have? Would Na give up its valence electron spontaneously? Why does Na tend to form an Na + cation, rather than an Na 2+ (or any other) cation? What about Mg? What about F?
Anion formation How many valence electrons does F have? Is it possible to add an extra electron into the valence energy level of F? Why does F tend to form an F - anion, rather than an F 2- (or any other) anion? What about O? What about Be?
Electron configurations Write the electron configuration for each neutral atom. What ion will form for each? F O Li He Ne Cl
Symbols for ions: Use the chemical symbol and the charge for each ion. Notice PT trends! Some common ions
Periodic relationship of simple ions Transition Metals?
Ionic bonding Oppositely charged ions attract! This forms a strong electrostatic attraction Called an Ionic Bond.
Atoms will become charged before making ionic bonds Ion charged atom loss or gain of electrons Cations: Anions: (fulfills octet or duet rule) Use periodic table to predict the ion formed for O, Li, He, Ne, and Cl
Ionic bonds-form a crystal lattice Ions form ionic crystals. pattern of alternating + and – ions An exothermic process. Crystal structure extremely stable. Click HERE for an animated tutorialHERE
Ionic Solids Salts are held together by ionic bonds
Ionic compounds When cations and anions come together to form ionic bonds, an ionic compound is the result. All compounds are neutral. Ex: Na reacts with S
Ionic Compounds Most ionic compounds are composed of a metal and a nonmetal. Example: Sodium Chloride Example: Magnesium Oxide Are ionic compounds charged? Ionic compounds are NEUTRAL (no charge), so the two ions together must cancel each other’s charge. Example: Mg +2 and O -2 or Na + and Cl - One of each ion, makes a neutral ratio.
Ionic Compounds What if the ions are of different charges? Example: Sodium and Oxygen Na +1 and O -2 The ions exist in ratios, producing an electrically neutral compound. 2 sodium ions for every oxygen ion How can we write this as a formula? Na 2 O This formula shows NO charges, it is neutral There are two sodium ions for every oxygen It represents a RATIO of ions.
Practice writing formulas for the following ionic compounds: potassium fluoride calcium bromide aluminum chloride sodium sulfide beryllium oxide
Formulas of ionic compounds potassium fluoride KF calcium bromide CaBr 2 aluminum chloride AlCl 3 sodium sulfide Na 2 S beryllium oxide BeO
Rules for naming ionic compounds Salts (ionic compounds) are named: Cation first, anion second. Cations are named the same as the neutral atom, i.e. sodium, aluminum… they almost always end in “-ium”. Anions are named with the ending “-ide”. i.e. fluorine fluoride; oxygen oxide Examples: sodium chloride, magnesium oxide
Ions of transition metals Some transition metals can form more than one ion Roman numerals indicate which it is. Example: Name: Nickel (I) oxide: Ions: Ni +1 and O -2 Formula:Ni 2 O Example: Name: Cobalt (II) Chloride Ions:Co +2 and Cl -1 Formula:CoCl 2
Roman numeral exceptions Some transition metals only form one stable ion and do not require roman numerals: Ag + Zn +2 A few other metals have more than one stable ion and require roman numerals: tin(II) and tin(IV) lead(II) and lead(IV) Memorize these...
Practice writing names for ionic compounds of transition metals ZnCl 2 PbBr 2 AuO
Names of ionic compounds ZnCl 2 zinc chloride PbBr 2 lead (II) bromide AuO gold (II) oxide
Polyatomic ions: to memorize Polyatomic ion Name: SO 4 -2 sulfate NO 3 -1 nitrate CO 3 -2 carbonate HCO 3 -1 bicarbonate NH 4 +1 ammonium OH -1 hydroxide PO 4 -3 phosphate CH3COO -1 (C 2 H 3 O 2 -1 ) acetate
Introducing dimensional analysis. Suppose you work at a hardware store and a customer comes in asking to buy 150 bolts from a bulk bin. What would be the most efficient way to get the customer his 150 bolts.
Counting large numbers Atoms are too small to count one at a time! The “mole” is a useful unit for counting atoms or other small particles. 1 mole = 6.022 x 10 23 objects
Compare and contrast Compare 5 dozen bagels to 5 dozen elephants How are they the same? How are they different? Compare 2.5 moles of hydrogen atoms to 2.5 moles of gold atoms How are they the same? How are they different?
Avogadro’s Number Ratio between grams and a.m.u. Therefore, atomic masses on PT are also the molar masses of elements (in g/mol)
Counting by weighing We know the mass of 1 mole of any element AND We know how many atoms are in 1 mole of any element THEREFORE we can “count” a certain number of atoms using mass.
Molar mass of compounds To calculate the mass of one mole of a compound, add the molar masses of the atoms EX: 1 mole of Al 2 O 3 has 2 mol Al + 3 mol O What is the molar mass of: Li 3 P NaOH Mg(NO 3 ) 2
Mole ratios for compounds One mole of CaCl 2 has mole(s) of Ca 2+ ions mole(s) of Cl - ions One mole of Al 2 (CO 3 ) 3 has mole(s) of Al 3+ ions mole(s) of CO 3 2- ions
Examples: What is the mass of 1 mole of calcium atoms? What is the mass of 2 moles of calcium atoms? What about 5.3 moles? How many atoms are there in 2.25 moles of sulfur? What is the mass of this sample?
Calculations with ionic compounds How many moles in 5000 kg of iron(III) oxide? What is the mass of 2.50 mol of calcium chloride? How many potassium ions are there in 4.5 moles of potassium oxide? How many formula units are there in 285 g of copper(II) oxide?
Test Preparation: Start the review sheet…but don’t end there. Use practice problems from the book at the end of the chapter. Power point DO PRACTICE PROBLEMS! Do the worksheets or look over labs again. Be able to explain ideas completely. Cause and effect.
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