A. ATOMS Name Period Date 1. Complete the following table. Element Symbol Number of Protons. Number of electrons Ac 227

Size: px

Start display at page:

Download "A. ATOMS Name Period Date 1. Complete the following table. Element Symbol Number of Protons. Number of electrons Ac 227"

Georgina Cross
5 years ago
Views:

1 A. ATOMS Name Period Date 1. Complete the following table. Element Symbol Number of Protons Number of electrons Number of neutrons Atomic Number Mass Number Ac Fill in the following Table Element Symbol Atomic Number Mass Number Number of neutrons nitrogen Ne Beryllium Use the following information to determine the atomic mass of chlorine. Two isotopes are known: chlorine-35 (mass = amu) and chlorine-37 (mass = amu). The relative abundance s are 75.4% and 24. 6%, respectively. 4. Use the following information to determine the atomic mass of carbon. Two isotopes are known: carbon-12 (mass = amu) and carbon-13 (mass = amu). Their relative abundance s are 98.9% and 1.10% respectively. 5. Given the relative abundance of the following naturally occurring isotopes of oxygen, calculate the average atomic mass of oxygen. Assume that the atomic mass of each is the same as the mass number. oxygen- 16: 99.76% oxygen- 17: 0.037% oxygen-18: 0.204% 6. Distinguish between protons, electrons, and neutrons in terms of their relative masses and charges. 7. Discuss the structure of an atom including the location of the proton, electron, and neutron with respect to the nucleus.

2 8. Summarize Dalton s atomic Theory 9. In what type of ratios do atoms combine to form compounds? (amu) atomic number periodic table mass number group isotopes atomic mass unit atomic mass period electrons cathode ray protons neutrons nucleus atom scanning tunneling electron microscope John Dalton Democritus A. atoms that have the same number of protons but different numbers of neutrons B. weighted average mass of the atoms in a naturally occurring sample of an element C. equals the number of neutrons plus the number of protons in an atom D. 1/12 the mass of a carbon-12 atom E. the number of protons in the nucleus of an atom of an element F. an arrangement of elements according to similarities in their properties G. a vertical column of elements in the periodic table H. a horizontal row of the periodic table I. stream of electrons produced at the negative electrode of a tube containing a gas at low pressure J. the central core of an atom, which is composed of protons and neutrons K. negatively charged subatomic particles L. subatomic particles with no charge M. positively charged subatomic particles N. an instrument used to generate images of individual atoms O. Greek philosopher who was among the first to suggest the existence of atoms P. the smallest particle of an element that retains its identity in a chemical reaction Q. English chemist and schoolteacher who formulated a theory to describe the structure and chemical reactivity of matter in terms of atoms

3 B. PERIODIC TABLE 1. Periodic means. Examples of periodic properties: 2. What is a group (or family)? What is a period? 3. How can you determine the number of electrons in an element s outer energy level by the group it s in? 4. What is the octet rule? 5. Why do elements that make positive ions occur on the left side of the periodic table while those that make negative ions occur on the right? 6. What is the common name for group 18? Why do the elements of this group usually not form ions? 7. Complete the following table. Group Common Name Charge on Ions of this Group / 3A / 6A / 7A 8. Predict the charges on ions of the following atoms. Ra As Te Cs In At Ga 9. a) In group 1, which element is the most active? b) Metallic activity tends to (increase, decrease) as one goes down Group a) Which element is most active in group 17? b) Nonmetal activity tends to (increase, decrease) as one goes down Group Compare and contrast ionization energy and atomic radius. Ionization Energy Radius Definition: Largest values (metal or nonmetal side) Largest values (top or bottom of group)

4 Use a periodic table to help you answer the following questions. 1. Which element in the second period has the greatest atomic radius? 2. Which of the group IIIA (13) elements is the largest? 3. Of the halogens, which has the smallest radius? 4. Which of the alkaline earth metals is the largest? 5. Which of the transition metals has the smallest atomic radius? 6. Which of the noble gases is the smallest? 7. The atomic radius of which element is the largest? 8. Do alkali metals generally make anions or cations? 9. Which of the elements which have their valence electrons in the second energy level is the largest? 10. Which of the metalloids has the smallest atomic radius? 11. Which of the rare earth elements is the smallest? 12. Which of the transition metals in the fifth period is the largest? 13. Are metal ions larger or smaller than the neutral atoms they came from? 14. Are cations larger or smaller than the neutral atoms they came from? 15. Are ions of alkali metals larger or smaller than ions of alkaline earth metals from the same period? 16. Which element in the second period has the greatest first ionization energy? 17. Which of the group IIIA (13) elements has the largest ionization energy? 18. Of the halogens, which has the smallest electronegativity? 19. Which of the alkaline earth metals has the smallest electronegativity? 20. Which of the transition metals has the largest ionization energy? 21. Which of the noble gases has the smallest ionization energy? 22. Which of the group IVB (14) metals is the least active? 23. Which of the halogens is the most active?

5 24. Which of the semi-metals that have their valence electrons in the fourth energy level has the largest ionization energy? 25. Which of the period three elements has the largest electronegativity? 26. Which of the inner transition elements of the seventh period is the easiest to ionize? 27. Which of the transition metals in the fifth period has the largest Electronegativity N? 28. Which of the group four metals has the largest ionization energy? 29. Which of the non-metals in the third period is the most active? 30. As atomic size increases, what happens to the ionization energy of the atom?

6 B. ELECTRONS IN ATOMS AND BONDING Determining the Ionic Charge Element Property Before Making an Octet After Making an Octet electron 1s 2 2s 1 1s 2 config # protons 3 3 #electrons 3 2 charge ZERO 1+ Li Bohr Diagram Lewis Dot Structure electron config # protons #electrons charge Be Bohr Diagram Lewis Dot Structure electron config # protons #electrons charge B Bohr Diagram Lewis Dot Structure

7 Element Lewis Dot Periodic table group # of Valance e- Gain/Lose e- Valance Charge Na Na. 1 1 LOSE 1 +1 Be Cl S Al Ne K N O Ca P B Mg

8 Write the Formula / Formula Unit for the following Compounds Determining the formula for Magnesium Fluoride? 1. Identify the charges = Mg 2+ F 1-2. Cross the Values of Charges, Mg 2+ F 1- = MgF 2. Notice: No +/- signs and the numbers are SUBSCRIPTS and because the subscript would be a 1, it does not need to be written. 3. If there is a common subscript such as 2 as in Mg 2 O 2, reduce it to MgO. Write Formula Unit For the Below Ionic Compounds Name Cation (+) Anion (-) Formula Sodium Chloride Na 1+ Cl 1- Na 1+ 1 Cl 1-1 = NaCl Lewis Dot Diagram Aluminum Chloride Aluminum Phosphide Al 3+ Cl 1- Magnesium Oxide Cesium Fluoride Strontium Nitride Lithium Sulfide Calcium Chloride Sodium Bromide Beryllium Iodide Strontium Fluoride Aluminum Fluoride Potassium Nitride Sodium Sulfide Lithium Oxide Calcium Oxide

9 C. NAMES AND FORMULAS Polyatomic Ions Ammonium (NH 4 ) 1+ Acetate (C 2 H 3 O 2 ) 1- Carbonate (CO 3 ) 2- Bicarbonate (HCO 3 ) 1- Chromate (CrO 4 ) 2- Chlorate (ClO 3 ) 1- Perchlorate (ClO 4 ) 1- Chlorite (ClO 2 ) 1- Hypochlorite (ClO) 1- Dichromate (Cr 2 O 7 ) 2- Hydroxide (OH) 1- Nitrate (NO 3 ) 1- Nitrite (NO 2 ) 1- Sulfate (SO 4 ) 2- Sulfite (SO 3 ) 2- Phosphate (PO 4 ) 3- Permanganate (MnO 4 ) 1- Hydrogen phosphate (HPO 4 ) 2- Write the NAME of each of the following compounds. 1. (NH 4 )Cl 2. Be(SO 4 ) 3. (NH 4 ) 3 N 4. MgCl 2 5. NH 4 (NO 3 ) 6. Sr 3 (PO 4 ) 2 7. Zn(CrO 4 ) 2 8. K 2 (Cr 2 O 7 ) 9. Ga(ClO 3 ) Cu(OH) 11. (NH 4 ) 3 (PO 4 ) 12. Fe(SO 4 ) 13. Mg(NO 3 ) (NH 4 )NO Na 2 (Cr 2 O 7 ) 16. Na(OH)

10 Write the CHEMICAL FORMULA for each of the given NAMES cross charges. 17. calcium carbonate Ca 2+ (CO 3 ) 2- = Ca 2 (CO 3 ) 2 = CaCO calcium nitrate = 19. ammonium sulfate = 20. aluminum hydroxide = 21. barium phosphate = 22. cesium nitrate = 23. sodium nitrite = 24. calcium sulfate = 25. beryllium sulfate = 26. sodium hydrogencarbonate = Write the formula for the variable charged binary ionic compounds: 1. Nickel (II) chloride 2. Gold (III) oxide 3. Cobalt (II) phosphide 4. Copper (I) chloride 5. Iron (III) chloride 6. Copper (II) bromide 7. Vanadium (V) sulfide 8. Cobalt (II) phosphide 9. Manganese (III) phosphate 10. Iron (III) oxide

11 NAME the Ionic Compound 19. Ca Br Fe O 21. Cu S 22. Cr N 23. V 3 (PO 4 ) Li 2 S 25. Mg(NO 3 ) Ni 3 (PO 4 ) 2 D. Ions in Chemical Compounds Complete the following table, being sure that the total charge on the resulting compound is zero. Ions Chloride Cl 1- Hydroxide (OH) 1- Nitrate (NO 3 ) 1- Sulfate (SO 4 ) 2- Sulfide S 2- Carbonate (CO 3 ) 2- Phosphate (PO 4 ) 3- Cyanide (CN) 1- Ammonium NH 4 1+ Potassium K Calcium Ca Magnesium Mg Aluminum Al Iron (II) Fe Strontium Sr Iron (III) Fe

12 Lead (II) Pb Tin (IV) Sn Copper (I) Cu Hydrogen H E. CHEMICAL BONDING Ionic Bond between a Metal and Non-Metal (M + NM) Covalent Bond between a Non-Metal and Non-Metal (NM + NM) Metallic Bond between a Metal and Metal (M+ M) Determine if the elements in the following compounds are metals or non-metals. Describe the type of bonding that occurs in the compound. Compound Element 1 Element 2 Bond Type (metal or non-metal?) (metal or non-metal?) NO 2 N = non-metal O = non-metal covalent NaCl SO 2 PO 4 3- MgBr 2 CaO H 2 O K 2 O Cu-Zn alloy O 2 CuCl 2 NO 2 - TiO 2 HF

Rb 2 S Hg-Ag amalgam Fe 2 O 3 C 6 H 12 O 22 Electronegativity: A property of an atom which increases with its tendency to attract the electrons of a bond.

13 Rb 2 S Hg-Ag amalgam Fe 2 O 3 C 6 H 12 O 22 Electronegativity: A property of an atom which increases with its tendency to attract the electrons of a bond. Examples: The chlorine atom has a higher electronegativity than the hydrogen atom, so the bonding electrons will be closer to the Cl than to the H in the HCl molecule. Electronegativity Difference in electronegativity Polar-covalent Ionic bond Non-polar covalent bond 100% 50% 5% 0% Percentage Ionic character Bonding between Sulfur and Hydrogen Sulfur and cesium Chlorine and bromine More electronegative element and value Less electronegative element and value Difference in electronegativity Bond Type

14 Calcium and chlorine Oxygen and hydrogen Nitrogen and hydrogen Iodine and iodine Copper and sulfur Hydrogen and fluorine Carbon and oxygen BALANCE THE FOLLOWING EQUATIONS TO USE IN QUESTIONS 5 through 14 below: 1. Al + O 2 Al 2 O 3 2. Cu + AgNO 3 Ag + Cu(NO 3 ) 2 3. Zn + HCl ZnCl 2 + H 2 4. Fe + Cl 2 FeCl 3 PERFORM THE FOLLOWING STOICHIOMETRIC CALCULATIONS: 5. Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen. How many moles of HCl are required to produce 7.50 moles of ZnCl 2? 6. Copper metal reacts with silver nitrate to form silver and copper(ii) nitrate. How many grams of copper are required to form 250 g of silver? 7. When aluminum is burned in excess oxygen, aluminum oxide is produced. How many grams of oxygen are required to produce 0.75 moles of Al 2 O 3? 8. How many grams of iron(iii) chloride are produced when 15.3 g of iron react with excess chlorine gas? 9. Copper metal reacts with silver nitrate to form silver and copper(ii) nitrate. How many moles of silver will be produced from 3.65 moles of silver nitrate? 10. Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen gas. How many milliters of 3.00M HCl are required to react with g of zinc? 11. How many grams of iron are needed to react with 31.0 L of chlorine gas at STP to produce iron(iii) chloride? 12. When 9.34 g of zinc react with excess hydrochloric acid how many grams of zinc chloride will be produced? 13. How many liters of oxygen gas at STP are required to react with 65.3 g of aluminum in the production of aluminum oxide? 14. Copper reacts with silver nitrate to form silver and copper(ii) nitrate. How many grams of copper are required to react with 50.0 ml of 8.0M AgNO 3

15 SOLVE THE FOLLOWING LIMITING REACTANT PROBLEMS: 15. When 16.3 g of magnesium and 4.52 g of oxygen gas react, how many grams of magnesium oxide will be formed? Identify the limiting and excess reactants. 2 Mg + O 2 2 MgO 16. If 25.3 g of aluminum reacts with 25.3 g of copper(ii) sulfate, how many grams of copper are formed? Identify the limiting and excess reactants in this single replacement reaction. 2 Al + 3 CuSO 4 3 Cu + Al 2 (SO 4 ) Identify the limiting and excess reactants when 1.00 g of zinc reacts with 150 ml of 0.250M Pb(NO 3 ) 2. How many grams of lead are formed in this single replacement reaction? Zn + Pb(NO 3 ) 2 Pb + Zn(NO 3 ) If 24.5 g of iron are placed in 1.00 L of 0.25M HCl, how many grams of FeCl 2 are obtained? Identify the limiting and excess reactants in this single replacement reaction. Fe + 2 HCl FeCl 2 + H 2 SOLVE THE FOLLOWING PERCENT YIELD CALCULATIONS: 19. If 12.5 g of copper react with excess chlorine gas, then 25.4 g of copper(ii) chloride are produced. Find the theoretical and percent yields. Cu + Cl 2 CuCl If 6.57 g of iron react with an excess of hydrochloric acid, HCl, then 11.2 g of iron(ii) chloride are obtained in addition to hydrogen gas. Find the theoretical and percent yields. Fe + 2 HCl FeCl 2 + H If 5.45 g of potassium chlorate are decomposed to form potassium chloride, 1.75 g of oxygen gas is also given off. Find the theoretical and percent yields. 1. Explain why water is a polar molecule? 2 KClO 3 2 KCl + 3 O 2 2. How do hydrogen bonds form? 3. What are the two parts of a solution?

16 4. What types of molecules dissolve easily in water? What types do not dissolve easily in water? 5. What is surface tension? 6. Why does ice float on water? 7. How many different ways of helping a solid dissolve in a liquid are there? 8. How does temperature change the solubility of a solid in a liquid? How is it different from the solubility of a gas? 9. What is concentration of a solution and how is it expressed? 10. What is molarity and how is it different from molality? 11. What happens to the concentration of a solution when more liquid is added? 12. According to the Arrhenius concept, an acid is a substance that. A) is capable of donating one or more H + B) causes an increase in the concentration of H + in aqueous solutions C) can accept a pair of electrons to form a coordinate covalent bond D) reacts with the solvent to form the cation formed by autoionization of that solvent E) tastes bitter 13. A Bronsted-Lowry base is defined as a substance that. A) increases [H + ] when placed in H 2 O B) decreases [H + ] when placed in H 2 O C) increases [OH - ] when placed in H 2 O D) acts as a proton acceptor E) acts as a proton donor 14. A Bronsted-Lowry acid is defined as a substance that. A) increases K a when placed in H 2 O B) decreases [H + ] when placed in H 2 O C) increases [OH - ] when placed in H 2 O D) acts as a proton acceptor E) acts as a proton donor 15. A substance that is capable of acting as both an acid and as a base is. A) autosomal B) conjugated C) amphoteric D) saturated E) miscible 16. According to the following reaction, which molecule is acting as an acid? H2O + H2SO4 H3O+ + HSO4 -

17 A) H 2 SO 4 B) H 2 0 C) H 3 O + - D) HSO 4 E) None of the above 17. According to the following reaction, which molecule is acting as base? H2O + H2SO4 H3O+ + HSO4 - A) H 2 SO 4 B) H 2 0 C) H 3 O + D) HSO 4 - E) None of the above 18. In basic solution,. A) [H 3 O + ] = [OH - ] B) H 3 O + ] > [OH - ] C) [H 3 O + ] < [OH - ] D) [H 3 O + ] = 0M E) [OH - ] > An aqueous solution contains 0.10 M NaOH. The solution is. A) very dilute B) highly colored C) basic D) neutral E) acidic 20. Nitric acid is a strong acid. This means that. A) aqueous solutions of HNO 3 contain equal concentrations of H + (aq) and OH - (aq) B) HNO 3 does not ionize at all when it is dissolved in water C) HNO 3 ionize completely to H + (aq) and NO - 3 (aq) when it dissolves in water D) HNO 3 produces a gaseous product when it is neutralized E) HNO 3 cannot be neutralized by a weak base 21. Which of the following solutions is the most acidic? A) a solution with ph = 3 B) a solution with ph = 5

18 C) a solution with ph = 7 D) a solution with ph = 10 E) all of the solutions are basic 22. If you had a 1.0 M solution of a strong acid, what would be a reasonable ph? A) 1 B) 6 C) 7 D) 8 E) If you had a 1.0 M solution of a weak acid, what would be a reasonable ph? A) 1 B) 6 C) 7 D)8 E) If you had a 1.0 M solution of a weak base, what would be a reasonable ph? A) 1 B) 6 C) 7 D) 8 E) If you had a 1.0 M solution of a strong base, what would be a reasonable ph? A) 1 B) 6 C) 7 D) 8 E) What are the H + and OH - concentrations that correspond to the following ph values? ph = 3.21 ph = 12.6 ph = 7.93 ph = 9.82 ph = 7.00 [H+] [OH-] 27. As the ph increases, the hydroxide ion concentration. A) decreases B) increases C) starts to affect the [H+] D) stays constant

19 28. If the ph of a solution is 10, what is the hydronium ion concentration? A) 4 B) 10 C) 1 x 10-4 D) 1 x 10-7 E) 1 x If the ph of a solution was 7 and you were to increase the hydroxide ion concentration, what would the ph be? A) 1 B) 5 C) 7 D) 9 E) 1 x 10-5 ===============================================================================. Calculation of the Final Equilibrium Temperature: For each of the following, draw a sketch, and solve the problem. 1. If grams of warm water at 60.0 o C is mixed with grams of cold water at 20.0 o C, find the equilibrium temperature. 2. If 25.0 grams of ice at 0.0 o C is added to grams of warm water at 40.0 o C, determine the final temperature. 3. Find the equilibrium temperature if grams of cold water at 10.0 o C is mixed with grams of warm water at 60.0 o C g of ice at 0.00 o C are added to 50.0 g of steam at 140 o C. Find T f. 5. A 600 ml sample of water is at 80.0 o C. How many grams of ice at o C must be added to bring the temperature down to 5.00 o C? 6. Butane has a M.P. = -125 o C, a B.P. = 35 o C, Specific Heat (solid) = 0.30 cal/g o C, Specific Heat (liq) = 0.70 cal/g o C, Specific Heat (gas) = 0.20 cal/g o C, H f = 50 cal/g, H v =?? 10.0 g of butane at -140 o C are added to 40.0 g of butane at 60.0C. T f = 0.0 o C. Find H v. B. Calorimetry Problems: 7. A 25.0 g sample of an alloy was heated to o C and dropped into a beaker containing 90 grams of water at o C. The temperature of the water rose to a final value of o C. Neglecting heat losses to the room and the heat capacity of the beaker itself, what is the specific heat of the alloy?

20 8. The specific of solid iron is cal/g o C. If 230 g of Fe at 95 o C were placed in 1 L of water at 15 o C, what would T f be? 9. Exactly 3 grams of carbon was burned to CO 2 in a copper calorimeter. The mass of the calorimeter is 1500 g, and the mass of water in the calorimeter is 2000 g. The initial temperature was 20.0 o C, and the final temperature is 31 o C. Calculate the heat value of carbon in calories per gram g of an unknown metal at 50 o C is placed in 40 g of water at 25.0 o C. The temperature of the system at equilibrium is 29.7 o C. What is the unknown metal? Your choices, and their Specific Heat values (in cal/g o C) are: a. Ba (0.068) b. Cr (0.111) c. Cu (0.092) d. Au (0.031) e. Sn (0.051) f. Rb (0.086) g. Ti (0.142) 11. A 10.0-gram hook, with a specific heat of 0.40 J/g o C, is heated to red-hotness in the flame of a Bunsen burner. The hook is quickly dropped into grams of water at 20.0 o C. If the temperature rises to 25.0 o C, then calculate the temperature of the Bunsen burner flame.

Ions and Ionic Compounds

Ions and Ionic Compounds Elements combine in a specific ratio to form compounds. Compounds can be categorized as ionic or covalent depending on the type of bond present within the compound. Ionic compounds