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Name__________________________ Date _____________ Period _______ Chemistry: Fall Final Review 08 Safety & Scientific Method 1) What are the safety rules for goggles, chemical safety, lab procedure, fire safety, diluting acids, etc. Review general safety rules; remember to dilute an acid always pour acid into water 2) Draw a picture of the biological, toxic, open flame, extreme temperature, and sharp object safety symbols. Biohazard Toxic Open Flame Extreme temp Sharp object 3) Give the location of safety devices such as the vent hood, eyewash, fire blanket, fire extinguisher, safety shower, etc. in the laboratory room. 4) List and describe the steps of the scientific method. Observations Hypothesis Experiment – will have a respondent (dependent) variable and an independent (manipulated) variable Analyze and conclude results Name__________________________ Date _____________ Period _______ 5) Define qualitative and quantitative. Qualitative – deals with the type of substances Quantitative – deals with the amount of substances Chapter 1 & 3 – Introduction to chemistry 6) Define chemistry. Chemistry – the study of the structure and composition of matter 7) Name and define the five types of chemistry. 5 types of chemistry: Organic chemistry – study of substances containing carbon Inorganic chemistry - study of substances not containing carbon Analytical chemistry – study of the composition of substances Physical chemistry – study of the physical behavior of chemicals Biochemistry – study of the chemistry of living things 8) Define mass and matter. Mass – the amount of matter in an object Matter – anything with mass and takes up space 9) Define and describe the states of matter. Solid – definite shape and volume Liquid – definite volume Gas – no definite shape or volume Plasma 10) What is a physical property? Give examples. Physical properties can be observed using your senses (color, melting point, state of matter, etc.) 11) What is a chemical change? Give examples. Chemical changes occur when something new is produced (burning, tarnishing, reacting, etc.) Name__________________________ Date _____________ Period _______ 12) What are the differences between elements and compounds? Elements are made of only one type of atom; Compounds are made of more than one type of atom 13) What is a mixture? Physical blend of 2 or more pure substances in any proportion in which each substance retains its individual properties 14) What are the differences between a homogeneous and heterogeneous mixture? Homogeneous mixture – you can’t see the different parts Heterogeneous mixture – you can see the different parts 15) What are the four evidences that a chemical reaction has taken place? Bubbles, heat, color change, precipitate 16) List & describe the separation techniques of mixtures. Distillation – separation of two liquids (based on boiling points) Filtration – separation of a solid and liquid Crystallization – forms a solid structure Chromatography – based on charge attraction to another material 17) Give the procedure for separating salt and sand and iron filings. Use a magnet to separate the iron; dissolve the salt in water and filter the sand 18) Define and give an example of the law of conservation of mass. 4 C3H5(NO3)3 4 CO2 + 6 N2 + 10 H2O + 9 O2 19) Define and give an example of the law of definite proportions? Regardless of the amount the elements will always be in the same ratio Ex. Water in a glass and water in a pool – H2O – ratio doesn’t change 20) Define and give an example of the law of multiple proportions? Two elements can combine in different ratios in order to give different compounds NO3 – nitrate Name__________________________ Date _____________ Period _______ NO2 – nitrite NO – hyponitrite (nitric oxide) 21) What is the difference between mass and weight. Mass – measurement of matter Weight – measurement of the force of gravity upon mass Chapter 2 – Dimensional Analysis & Measurement 22) What would 38 000 000 000 be in scientific notation? 3.8 x 1010 23) Subtract 1.2x107 – 9.1x106. 2.9 x 106 24) (1x103) x (5x10-1) = 5 x 102 25) Rules for Significant figures: 26) Know the rules for significant figures. How many are in: 1.70 Significant figures = 3 0.00550 3 20,000 1 10900 3 27) Define precision and accuracy. Give an example of a measurement that is precise but not accurate. Precise measurements are close to each other Accurate measurements are close to the actual value 28) % error – question – removed!!! 29) How many liters are in 12.7 kL? 12,700 Liters 30) Convert 22.5 L to mL. 22,500 mL Name__________________________ Date _____________ Period _______ 31) Use dimensional analysis to convert 22.5 km/hr into m/sec. 22.5 km x 1000 m x 1 hour x 1 min = 6.25 m/s 1 hour 1 km 60 min 60 sec 32) Give an example of everyday liquids (water, oil, Rubbing [isopropyl] alcohol) with different densities. How would they be arranged if poured together? Water is more dense than oil, therefore oil floats on top of water 33) What is the density in g/ml of a substance that has a mass of 3.99g and a volume of 0.88ml? D=m V D = 3.99 g D = 4.53 g/mL 0.88 mL 34) What is the volume of a liquid that has a density of 2.25g/ml and a mass of 18.7g? V=m D V = 18.7 g V = 8.31 mL 2.25 g/mL 35) What are the temperature conversions from Celsius to Kelvin? °C + 273 = K Chapter 4 – Atomic Structure 36) What was Democritus’ contribution to the atomic structure? He came up with the idea of atoms. 37) Give all of Dalton’s laws. Elements are composed of tiny indivisible particles called atoms; all atoms of the same element are identical; atoms can combine in whole number ratios to form compounds 38) Describe Rutherford’s gold foil experiment and what it was used to determine. Rutherford shot particles through a thin piece of gold foil; some of these particles occasionally bounced back giving Rutherford the idea that they were bouncing off of a dense area that he called the nucleus Name__________________________ Date _____________ Period _______ 39) What are the charges and locations of a proton, neutron and electron? Proton – positive Neutron – neutral Electron - negative 40) What information does the atomic number give you? The atomic number tells you the number of protons or the number of electrons in a neutral atom 41) What is an isotope? Isotopes of an element have a different mass number and a different number of neutrons 42) Define atomic mass. The weighted average mass of the isotopes of that element 43) Define mass number. Number of neutrons + protons 44) How do you find the number of neutrons in an atom? How many neutrons are in carbon-14? 14 – 6 = 8 neutrons in Carbon 14 # neutrons = mass # - atomic # 45) What are the different atomic models? Which is the most current. Plum Pudding Model – only showed electrons Rutherford’s Model – included a nucleus Bohr Model – showed electrons in energy levels Quantum Mechanical Model – shows the probable locations of electrons using the solutions from Schrodinger’s equations 46) Define and give an example of Alpha, Beta, and Gamma decay. 1. Alpha (α) – an alpha particle is 2 protons and 2 neutrons (42α) 206 4 Ex. 21084Po 82Pb + 2α 2. Beta (β) – a beta particle is the lost of an electron (when a neutron is converted to a proton) ( 0-1e-) 14 0 Ex. 146C 7N + -1e3. Gamma (γ) – a gamma particle is a photon at high frequency 31 Ex. 3115P* 15P + γ Chapter 5 - Electrons 47) What formula for the speed of light? c = λƒ Name__________________________ Date _____________ Period _______ 48) What is the formula for the Energy of a photon? E = hƒ 49) Define photon. A particle of electromagnetic radiation with no mass that carries a quantum of energy 50) Define atomic emission spectra. What does it tell us about the atoms of elements? A set of frequencies of electromagnetic waves given off by atoms of an element 51) Define Heisenberg uncertainty principle. It is not possible to know precisely both the velocity and the position of a particle at the same time. 52) Define Aufbau’s principle. Electrons enter orbitals of the lowest energy first. 53) Define Pauli exclusion principle. No 2 electrons in the same atom can have the same set of 4 quantum numbers 54) Define Hund’s Rule. Electrons in the same subshell occupy available orbitals singly before pairing up (1 e- per orbital till each orbital has one then begin putting a 2nd e- in each orbital) 55) Define Valence electrons. The electrons in the outermost orbital Help determine an atoms properties 56) Define Octet rule. Atoms lose or gain or share electrons in order to aquire a full set of 8 valence electrons. (first shell has 2 valence electrons but all others have 8) 57) How many electrons can each orbital hold? How many electrons can the s, p, d and f sublevels hold? Each orbital can only hold 2 electrons S sublevel – has one orbital = 2 electrons total P sublevel – has 3 orbitals = 6 electrons total D sublevel – has 5 orbitals = 10 electrons total 58) How can you easily locate the energy level by using periodic table? Locate the row of the element Name__________________________ Date _____________ Period _______ 59) Draw the orbital notation for chlorine. 60) What is the electron configuration of lead. 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2 61) In 2s1, what do the 2, the s and the 1 represent? The 2 represents the energy level and the 1 represents one electron. 62) How many valence electrons are in N, Cl and Na? N has 5 valence e- , Cl has 7 valence e- , Na has 1 valence e- (same as the group number) 63) What are the 4 quantum numbers of the last electron in Cr, S, I, Mg? OMIT!!! Cr – (3d4) - 3,2,1, ½ S – (3p4) – 3,1,-1, - ½ I – (5p5) – 5,1,0, - ½ Mg – (3s2) – 3,0,0, - ½ 64) Write out a stable electron configuration of Copper. [Ar] 4s23d9 - normal [Ar] 4s13d10 - [Ar] 4s03d10 – CU+1 [Ar] 4s23d9 - [Ar] 4s03d9 – CU+2 Chapter 6 – Periodic Table 65) How was Mendeleev’s periodic table arranged? By increasing atomic mass 66) How was Mosley’s periodic table arranged? By increasing atomic number 67) What is the group 1A elements called? Group 2A? Group 7A? Group 8A? What kind of elements are in groups 3-12? Group 1A – Alkali Metals Group 2A – Alkaline Earth Metals Group 17 (Group 7 A) – Halogens Group 18 (Group 8 A) – Noble Gases Groups 3-12 – Transition Elements Name__________________________ Date _____________ Period _______ 68) How are elements grouped on the periodic table? Valence electrons 69) What is the trend for atomic radius? What element has the largest atomic radius? Atomic radius increases from top to bottom on the periodic table and it increases from right to left. The largest atomic radius is Francium 70) Place in increasing radius - carbon, a carbon cation or a carbon anion? Carbon cation < Carbon atom < carbon anion 71) What is ionization energy? What is it’s trend? Ionization energy is the amount of energy it takes to remove an electron from an atom. Ionization energy increases from bottom to top and it also increases from left to right on the periodic table. Helium has the highest ionization energy. 72) What is the trend in electronegativity? Where are the elements with the smallest electronegativity? Electronegativity increases from left to right (not including the noble gases) and it increases from top to bottom. The smallest electronegativity elements are the noble gases; the biggest is Fluorine. Chapter 8 – Ionic Bonding 73) Define Ionic Bond. Complete transfer of electrons from one atom to another Bond between ions (cation and anion) 74) Define cation and anion. Cation – a positive ion resulting from the loss of electrons Anion – a negative ion resulting from the gain of electrons 75) Gaining three electrons gives what overall charge? Negative 3 charge Name__________________________ Date _____________ Period _______ 76) Draw the Lewis dot structure for N, K, & Cl-. N K Cl- 77) Name the following compounds: CaCl2 - Calcium Chloride CuNO3 - Copper (I) Nitrate CCl4 - Carbon Tetrachloride Ba(C2H3O2)2 - Barium Acetate Si3H8. - Trisilicon Octahydride 78) Write the formula for the following: zinc bromide ZnBr dinitrogen monoxide - N2O lead (IV) oxide - Pb2O copper (I) carbonate - Cu2CO3 calcium fluoride. – CaF2 79) When is an ionic compound able to conduct electricity and why? An ionic compound can conduct electricity when it is in solution (in water). This allows the ions to float around freely and conduct current. 80) What properties of metals are explained by metallic bonding theory? Metals are ductile and malleable due to their metallic bonding and the sea of free floating electrons. 81) Define and give an example of an alloy. Mixture of 2 metals. Steel Name__________________________ Date _____________ Period _______ Chapter 9 – Covalent Bonding 82) Give examples of pairs of elements that would form ionic bonds, nonpolar bonds and polar bonds. Nonpolar bonds are between elements with less than a 0.4 difference in electronegativity (ex: Cl2) Polar bonds have elements with between a 0.4 and 1.7 difference in electronegativity (ex: CCl4) Ionic bonds have more than a 1.7 difference in electronegativity (NaCl) 83) What are sigma and pi bonds? Sigma bond – formed by end to end overlap of an s orbital Pi bond – formed by side to side overlap of p orbitals 84) Compare and contrast ionic compounds to covalent compounds? Ionic compounds transfer electrons between atoms. Covalent compounds share electrons. 85) What is VSEPR? Valence Shell Electron Pair Repulsion Electron move as far away from each other as possible 86) How do you determine Molecular shape of a molecule? By the shared and unshared valence electrons (atoms attached and the lone pairs) 87) Draw the Lewis structure and list the molecular shape for these molecules: SeF2 - Bent CO2 - Linear CH4 - Tetrahedral NH3 – Trigonal pyramidal HCl - Linear CH3CH2OH – central atom - tetrahedral