Star Test Review

1. Star Test Review CHEMISTRY

2. BREAKDOWN OF QUESTIONS

3. 6 Questions on the Periodic Table • To relate the position of an element in the periodic table to its atomic number and atomic mass • Use the periodic table to identify metals, semimetals, non-metals, and halogens • Identify alkali metals, alkaline earth metals and transition metals, trends in ionization energy, electronegativity, and the relative sizes of ions and atoms…(2 Questions on this…) • Use the periodic table to determine the number of electrons available for bonding • The nucleus of the atom is much smaller than the atom yet contains most of its mass.

4. The Modern Model Dalton’s atom Thompson’s electrons Rutherford’s space and nucleus Bohr’s energy levels Chadwick’s neutrons (not to scale)

5. Elements 8 O • 110 known elements • 92 of which are naturally occurring. • Each has an atomic symbol. • Atomic number • is number of protons • Atomic mass • is the total mass of the protons plus the neutrons. OXYGEN 15.9994 Notice that the atomic mass is not a round number, even though protons and neutrons each have a mass of 1. This is due to natural abundance.

6. Energy Levels (n) • The electrons exist in energy levels or shells. • The first energy shell can hold only 2 electrons. • Hydrogen and Helium in their ground state have electrons that occupy this shell. • The second shell can hold 8 electrons. • The third can hold 18 electrons. 2 8 18 32 Shells All shells after three can hold 32 electrons.

7. Ion e- configurations • Ions(elements with more/less electrons) also have electron configurations. • Consider Sulfur (S): • What if sulfur gained an electron? • Consider Calcium (Ca): • What if calcium lost two electrons?

8. Question: • Why do the atomic radii of atoms decrease as electrons are added to the atom, as you move from left to right across a period? • electrostatic attraction • attraction between the electrons (-) in the shells and the protons(+) in the nucleus – pulls the electrons in This is what we call a periodic trend

9. The Periodic Table • The Periodic Table • a collection of all the known elements into a model that groups elements with similar properties. • Groups • Vertical columns of elements with similar properties. • Periods • Horizontal rows of elements with atomic mass and similar electron configurations.

10. Periodic Table History c 1869 • Dmitri Mendeleev • Russian chemist created who ordered the known elements according to properties. (Gaps?) • Henry Moseley • arranged the elements according to atomic number(# of protons). • This is the system we use today. • Periodic Law • chemical and physical properties of elements are periodic functions of their atomic numbers. • The elements in the periodic table are arranged according to Periodic Law • Periodic Law shows certain trends in the properties of elements c 1911

11. Periodic Trends – Atomic Radii • As electrons are added to the outside of atoms, in the same period, the atom’s radius decreases. • As new shells are added, radius increases. Smaller from left to right

12. Periodic Trends – Ionization Energy • Ionization Energy - the energy required to strip an electron from an atom. • As more electrons are added to a shell, it’s more difficult to remove them. (More protons pulling inward) • Easier to remove electrons from larger atoms. Larger from left to right

13. Period Trends – Electronegativity • Electronegativity (electron affinity) • an atom’s ability to attract electrons • Negative electron affinity = atom wants e-. • Decreases down a group Larger from left to right

14. Ionic Radii • Recall:+-attraction determines the atom’s radius. • An electron is added to a nonmetal atom : • Anion is formed. • Anions are larger than their neutral atom • An electrons is removed from a metal atom: • Cation is formed. • Cations are smaller than their neutral atoms Cl Cl- Na Na+

15. Groups and their Properties • Recall: • elements in the same group have similar properties due to similar electron configurations. • Learn the following group-families and their basic chemical and physical properties: • Alkali Metals • Alkaline-Earth Metals • Transition Metals • Main-Block Elements • Noble Gasses • Rare-Earth Elements

16. Group 1 (+1)Alkali Metals (s) • soft, highly reactive metals. • Lustrous • will reflect light, but these elements quickly lose their sheen when exposed to the air. • Electrically Conductive • able to pass a charge through the material. These elements are often found in lights, batteries, and electrolytes. • have low melting points • low density.

17. Group 2 (+2)Alkaline-Earth Metals (s) • Properties are similar to group 1 elements, but are: • Harder • Less reactive than Group 1 elements. • (These elements are still very reactive.) • Lustrous • Electrically Conductive • Higher melting points than Group 1 metals. • More dense than Group 1 metals.

18. Groups 3-12 (various)Transition Metals (d) • This is where we find most metals, including the coinage metals. • Lustrous • Electrically Conductive • Malleable • able to be shaped and formed, and hold that shape. • Ductile • able to be drawn into wires • Very hard • Verydense • High melting points

19. Group 13-17 (+3-1)Main-Block Elements (p) • The most varied elements. • Liquids, gasses, and solids can be found in this group. With widely varied properties • Includes Metalloids • elements having properties of both metals and non metals. • Most elements necessary to living things are found in this section. • Includes Halogens • Group 17 gasses and liquids F, Cl, Br, I, At • are very reactive due to very high electron affinities.

20. Group 18 (0)Noble Gases (p) • Nearly unreactive. • All have filled octets. • Near zero electron affinity • Very high ionization energies. • Noble gasses make up a trace amount of our atmosphere • are mined from pockets of gases in the oceans. • When electrically charged: • noble gases produce brilliant plasmas, often used in signs.

21. f – Group (various)Rare earth metals (f) • Very heavy, dense (large nuclei) • Most are radioactive. • Lanthanides • The first row, starting with lanthanum (57La) • (4f elements) • Actinides • The second row, starting with actinium (89Ac) • (5f elements) • Transuranium elements • All elements after Uranium92U (93Np on) are artificial.

22. Chemical Bonds---- 7 Questions • Atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds. (2 Questions) • Know chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2 and many large biological molecules are covalent. • Salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction • Atoms and molecules in liquids move in a random pattern relative to one another because the intermolecular forces are too weak to hold the atoms or molecules in a solid form • Know how to draw Lewis dot structures. (2 Questions)

23. Ionic bond: Electrons are given or taken, occurs between metal and nonmetal (NaCl) Chemical Bonding and Molecular Compounds Covalent bond: Electrons are shared, occurs between two or more nonmetals (CO2)

24. Ions +e- • Ion • atom that has gained or lost one or more electrons. • Octet Rule • all atoms want zero or eight electrons in their outer shell. • Metal and nonmetal bond: • each atom seeks to gain electrons or lose them. For instance: • Cl wants to gain one electron • Na wants to lose one electron • Notice the sizes of the ions Cl- Cl -e- Na Na+ Positive ions are called “cations” Negative ions are called “anions”

25. Ionic attraction Cl- Na+ • Oppositely charged ions attract • Like opposite poles of a magnet attract • Positively charged Sodium Ion and negatively charged Chlorine Ion are attracted. • They remain ions, but stick together in a lattice • (3D grid pattern) as other ions join them. Cl- Na+ Na+ Cl- Na+ Cl-

26. Formula Units • Na+ and Cl- combine to form NaCl. • NaCl is one formula unit: • the smallest unit that has the correct formula for our compound. • NaCl is also our compound’s empirical formula: • the smallest ratio of atoms that make up our compound. • Although there are several Na+ ions and Cl-ions in the lattice, the formula unit is still just NaCl. Cl- Na+

27. Metallic Bonding • In metallic bonding • electrons are shared and flow between metal atoms in an “electron sea.” • This is an extremely strong bond • This is why metals: • Have luster • (reflect light) • Are malleable • (able to formed and hold that form) • Are ductile • (able to formed into wires) • Conduct electricity • (pass an electrical charge) • Conduct heat • (pass heat easily) d-orbital electrons

28. Covalent Bonding • Covalent bonds are bonds formed when • atoms share electrons to complete their octet. • This is a very strong bond. • Atoms on the upper-right of the periodic table (nonmetals) covalently bond with each other. (except nobles) • Molecular compounds • covalently bonded compounds. • Diatomic molecules are covalently bonded. • I2, O2, F2, Cl2, H2, Br2, N2.

29. Covalent Bond Energy • Bond length and bond energy are inversely related. • That is, a smaller bond has a greater energy, and will release more energy when broken. • Atoms will bond at a distance that is most stable • (lowest potential energy) High potential energy = unstable Low potential energy = more stable

30. Electronegativity Cl- Na+ • Electronegativity: • Atom’s pull for electrons in a covalent bond • Electrons are sometimesnot shared equally among atoms. • Polar covalent • Electrons are shared, but one atom has a higher Electronegativity • Like in CO. • Nonpolar covalent • Electrons are shared evenly • Like in O2. C O O O

31. Molecular notation - Lewis Dot • Lewis Dot structures • tool used to draw molecules in 2D with stick diagrams. • Steps: • Determine each atom’s number of valence electrons • Number of valence electrons = number of dots • Molecules share dots until each atoms is surrounded by 8 dots, representing a completed octet. • Hydrogen gets only 2 dots on one side • Pairs of dots between atoms are changed to sticks – representing covalent bonds. • Any pairs of dots that are not shared between atoms are called unshared pairs.

32. Lewis Dot  Molecular Model • Consider Methane, CH4 • The atoms are C and H • Add each atom’s valence electrons as dots, clockwise.

33. Lewis Dot  Molecular Model • Consider Methane, CH4 • The atoms are C and H • Add each atom’s valence electrons as dots, clockwise. • Combine the dots to create octets, carbon belongs in the center.

34. Lewis Dot  Molecular Model • Consider Methane, CH4 • The atoms are C and H • Add each atom’s valence electrons as dots, clockwise. • Combine the dots to create octets, carbon belongs in the center. • Replace pairs of dots with sticks, representing bonds.

35. Lewis-Dot Practice: • Try the following on your own: • Determine each atom’s number of valence electrons • Arrange until all atoms except H have 8 valence electrons. (H has 2) • Replace ( : ) with sticks • (hint-carbon is always in the middle)

36. Conservation of Matter and Stoichiometry ---- 10 Questions • Know how to describe chemical reactions by writing balanced equations. ( 2 Questions) • Know the quantity one mole is set by defining one mole of carbon 12 atoms to have a mass of exactly 12 grams. • know one mole equals 6.02 x 1023 particles (atoms or molecules). • know how to determine the molar mass of a molecule from its chemical formula and a table of atomic masses and how to convert the mass of a molecular substance to moles, number of particles, or volume of gas at standard temperature and pressure. ( 3 Questions) • Know how to calculate the masses of reactants and products in a chemical reaction from the mass of one of the reactants or products and the relevant atomic masses. (3 Questions)

37. Chemical ReactionsProducts and Reactants • We write the compounds to react on the left and the compounds produced on the right. • Remember, reactants react to produce products Read as “yields”

38. Symbols and Notation • (s) – compound is a solid • (l) – cmpd is a liquid • (g) – gas • (aq) – aqueous (dissolved in water, exists as ions) • ↓ - a precipitate has formed • (solid falling out of the reaction) • ↑ - a gas is evolved • (bubbling out of the reaction) • Various things can be put over the “yields” sign (→) to indicate special reaction conditions.

39. Chemical Equations • We express a chemical reaction with a chemical equation. • This shows relative number of products and reactants required to satisfy the Law of Conservation of Mass. • Our lab reaction: • Must be written: • This balances the equation – both the reactants and products are equal.

40. Diatomic Molecules • Some elements don’t exist alone, but must form a pair with itself. • These are diatomic molecules. • In chemical reaction, these elements must be written as a X2. Diatomic Molecules Hydrogen = H2 Nitrogen = N2 Oxygen = O2 Fluorine = F2 Chlorine = Cl2 Bromine = Br2 Iodine = I2

41. Rules for Balancing Chemical Rxns • Never Change Subscripts • Balance Groups First • Balance H2O, O2, and H2 last. • We balance by adding coefficients. • Two reactant oxygens • Four product hydrogens • Properly balanced

42. Mole Ratios • Mole ratios: • how many moles of products are produced with given a number of moles of reactants. • Here, the mole ratio is 1:1:2 (1H2 : 1Cl2 : 2 HCl) • This means two moles of HCl will be produced when one mole of H2 and one mole of Cl2 react.

43. Balancing Practice 4 3 2 • Balance the following reactions on your own: Rust 2 2 Thermite + Energy 2 5 4 2 + Energy Acetylene torch 6 6 6 + Energy Respiration

44. Energy In and Out • Recall that a reaction can be Exothermic (releasing energy) or • Endothermic (absorbing energy) • Some exothermic reactions need a little energy to get going, but once going, will give off more energy. (energy barrier) • Activation Energy • energy needed to get the reaction going. Exothermic

45. The Mole • The “mole” represents a number of things….like a dozen. • How many things is a mole? • 6.022137 x 1023…but we will use 6.02 x1023. • This is Avogadro’s number • named for a lawyer, Amadoe Avogadro, that studied molecular gasses (diatomic) as a hobby. • When you have three moles of atoms, you have (3 x 6.02x1023 =) 1.81x1024 atoms total.

46. Molar Mass • Molar mass • expressed in grams per mole (g/mol) is the mass of one mole of a substance. • The mole is the link between the very small (atoms) and the macro (grams) • The average atomic mass of carbon is 12.01. What is the mass of a mole of carbon atoms? • Sodium has an atomic mass of 23 g/mol. How many moles do you have in 115 grams? 12.01 grams! 5.0 moles!

47. Mole-Mass Conversions • Complete the following mole-to-mass conversions: • Mass in grams of 2.25 moles of iron, Fe? • 126 grams Fe • Mass in grams of 0.375 moles of potassium, K? • 14.7 grams K • Number of moles in 5.00 grams of calcium, Ca? • 0.125 moles Ca • Number of moles in 3.60x10-10 grams of gold, Au? • 1.83x10-12 mol Au Use your periodic table to find molar mass

48. Phew! Mole-Atoms conversions • Mole = 6.02x1023 things, how many atoms are in: • 3.0 moles of silver, Ag? • 0.010 moles of copper, Cu? • How many moles do you have in: • 2.4x1024 atoms of helium, He? • 3.0x1023 atoms of lithium, Li? • How many moles do you have in 127.1 grams of copper? • How many atoms in 127.1 grams of copper?

49. Gases and Their Properties --- 6 Questions • Know the random motion of molecules and their collisions with a surface create the observable pressure on that surface. • Students know the random motion of molecules explains the diffusion of gases. • Know how to apply the gas laws to relations between the pressure, temperature, and volume of any amount of an ideal gas or any mixture of ideal gases. ( 2 Questions) • Know the values and meanings of standard temperature and pressure (STP). • Know there is no temperature lower than 0 Kelvin. • Know how to convert between the Celsius and Kelvin temperature scales.

50. Properties of Gases About 1000 times less dense than solids! • Gases have very low density • particles are spaced far apart. • Gases are compressible. • Extreme pressures-gases will compress until they become liquids (or solids, CO2). • At 1 atm and 0oC temperature (STP), • most gases have a standardmolar volume of about 22.4 liters. • One mole of gas at STP has a volume of 22.4 liters.