Benchmark 2 Review

1. Benchmark 2 Review

2. Percent Composition Percent Composition – percentage by mass of each element in the cmpd From sample amts: Mass of element in sample X 100=% E in compound mass of sample

3. Properties Chemists use properties to distinguish between substances and to separate them. By comparing several properties of a subst, an unknown substancecan be identified.

4. Properties Extensive property- depend on the amount matter such as mass Intensive property- does not depend on the amount of matter, such as odor.

5. Physical Properties • Physical Properties- can be observed or measured without changing the identity of the subst • Ex. Observing and measuring the boiling pt of water • Physical Change- a change that doesn’t involve changingthe identity of the subst • Ex. Liquid water changing to water vapor when boiled

6. Physical States of Matter Solid- definite shape and volume. Particles tightly packed in fixed positions and can only vibrate.

7. Liquid- definite volume, indefinite shape. Particles can flow past each other (fluid)

8. Gas- indefinite shape, indefinite volume, fluid. Particles move more rapidly with lots of space in between. Particles take up volume and shape of container

9. Plasma- high temperature matter where atoms lose their electrons. Ex. The sun

10. Chemical Properties • Chemical Properties- substance's ability to undergo changes to form new substances • Ex: ability of iron to react with oxygen to form rust

11. Periodic Table • Groups or Families - vertical columns of elements in the periodic table. • Numbered 1-18 OR with A/B system • Periods - horizontal rows of elements in the periodic table. • Numbered 1-7 Ex. Ge is located in group _____ and period _____

12. Metals and Nonmetals The periodic table is divided into two main sections: Metals and Nonmetals

13. Metals – • elements that are: Ex. Cu, Ni, Al • good conductors of electricity and heat • solids at room temp with luster • ductile (drawn into wires) • malleable (hammered into sheets) • tensile strength (resist breaking when pulled)

14. Nonmetals • elements that are : • Poor conductors of heat and electricity • Many gases • Brittle • Ex. C, O, S

15. Metalloids • Elements with characteristics of both metals and nonmetals • Ex. Si conducts electricity at high temp, not low temp; used as a semiconductor for computers

16. Noble Gases Noble Gases- elements in Grp 18 (8A) that are generally nonreactive Ex. He, Ne, Ar Section Review p. 24

17. Dmitri Mendeleev Dmitri Mendeleev developed the periodic table in 1869 1. He grouped elements with similar chemical and physical properties together 2. He arranged the elements in increasing atomic mass. 3. He predicted the properties of elements not yet discovered.

18. Periodic Table Today’s periodic table, the elements are arranged in increasing atomic number so elements with similar properties fall in the same group. Note: Ar and K both in according to atomic number, not atomic mass

19. Periodic Law Periodic Law: When elements are arranged in increasing atomic numbers, there is a repeating (periodic) pattern to the properties. See tables p. 142, 144, 148, 152.

20. Group Names • Group 1 elements – Alkali Metals – • Very, very reactive metals • Group 2 elements – Alkaline Earth Metals – • Very reactive metals • Group 17 elements – Halogens – • Very reactive nonmetals • Group 18 elements – Noble gases – • Nonreactive elements

21. Transition Metals • Group 3 to 12 – Transition Metals • Typical metals • Inner Transition Elements : lower two rows detached from main table • Lanthanides – atomic numbers from 58 to 71 • Actinides – atomic numbers from 90 to 103

22. Blocks of Elements S-block elements- valence electrons are filling the s orbitals- Groups 1-2 D-block elements-valence electrons are filling the d orbtials- Groups 3-12 P-block elements-valence electrons are filling the p orbitals- Groups 13-18 except He F-block elements-valence electrons are filling the f orbitals- Lanthanide and Actinide

23. Periodic Trends • Atomic Radii – ½ the distance between the nuclei of 2 of the same atoms bonded together. • Atomic radii trends: • Going across a period – atoms get smaller – caused by increasing positive charge. • Going down a group – atoms get larger – adding energy level Prac P.142

24. Ions Ion is an atom with a positive or negative charge. Cation formed by loss of electron(s) and has a positive charge. Anion formed by gain of electron(s) and has a negative charge. Illustrate example of atom, cation, and anion.

25. Ionic Radii • Ionic radii – positive ions are smaller than neutral atom due to loss of electron – negative ions are larger than neutral atom due to gain of electron • Ionic radii trends: same as atomic radii patterns • Going across a period – atoms get smaller – • Going down a group – atoms get larger –

26. Ionization Energy • Ionization energy (IE) – energy it takes to remove an electron from a neutral atom • IE trends: • Going across period – IE increases – greater positive to negative attraction, harder to remove electron • Going down a group – IE decreases – easier to remove an electron because it is farther from the nucleus

27. Electronegativity • Electronegativity – measure of the ability of an atom to attract electrons to itself • Electronegativity trends: • Going across a period – increase – more positive charge in nucleus in same energy level • Going down a group – decrease – nuclear attraction decrease with distance from nucleus Prac P.152 and section review p.154

28. Formulas • Empirical Formula – smallest whole number ratio of atoms in a compd. • Ex. HCl, H2O • Molecular Formula – actual formula of a molecule of compd – some multiple of the empirical formula • Ex. HCl, C2H4

29. Determining Empirical Formulas from Quantitative Analysis: Convert % info to grams. Convert grams to moles of element Get whole number ratios by dividing by smallest # of moles.

30. Molecular Formula Calculation of Molecular Formula: x(EF) = MF MF mass = x EF mass Problem must provide the molar mass of the compd from experimental data.

31. Measurements and Calculations Chapter 2

32. Units of Measurements A quantity must have magnitude, size, or amount. All scientist agree to use the same measurement system – the SI system.

33. SI Base Units – memorize these p34 Quantity, symbol Unit, abbrev. Length, l meter, m Mass, m kilogram, kg Temperature, T Kelvin, K Amt of substance, n mole, mol Time, t second, s

34. SI Prefixes – memorize these pg35

36. Derived units • Derived units – units obtained from a calculation • Area, A, m2 derived by: l x w

37. Volume, V, m3, derived by: l x w x h. We will also use non-SI volume unit of liter, L Memorize: one liter = one decimeter cubed (1 L = 1 dm3) one milliliter = one centimeter cubed (1 mL=1 cm3)

38. Derived units Density, D, kg/m3, derived by mass divided by volume D = M V Prob: 1. A 8.4 g sample of aluminum metal has a volume of 3.1 cm3. Calculate the density of aluminum.

39. Density Determine the volume of a 76.4 g sample of liquid mercury that has a density of 13.6 g/mL? Prac p.40; section review p.42

40. Conversion Factor Conversion factor is a ratio of equivalents of different units. Every conversion factor has a value of 1. Ex. 12 inches10-6 m 1 foot 1 um

41. Conversion Factor Prob. Write these into a conversion factor. 12 grams of copper are needed for every two students. A conversion factor between mg and g. A conversion factor between kL and L.

42. Dimensional Analysis Dimensional analysis – solving problems for the correct units as well as numbers. It is a way to help you set up the correct equation to solve the problem. Steps to solve dimensional analysis problem Write given on left. Set conversion factor so given unit will cancel. Base unit gets the number 1. Solve.

43. Scientific Measurements Accuracy – closeness to the accepted value or correct value Precision – closeness of a set of measurements to each other Look at dart examples p.44

44. Illustrate your own example of 5 quiz grades that have: High accuracy, high precision Low accuracy, high precision Low accuracy, low precision

45. Percent Error Percent error = valueaccepted – valueexperimental * 100 valueaccepted Problem: A student measures the mass of an object to be 4.80 g and the volume to be 3.2 mL. The actual density is 1.36 g/mL. What is the percent error of the student’s measurements? Practice p.45

46. Significant Figures See fig 2-9 on p46. Which value would be appropriate to record for the measurement? 6.3623678945678564 or 6.36 cm? Measurements in science are recorded in terms of significant figures to show the amount of certainty we have with the numbers that are being measured.

47. Significant Figures Significant Figures – a measurement with known digits and one estimated digit Rules to determine sign. fig. Non-zero numbers are significant. Count all zeros between nonzero numbers. Ex. 40.7 L has 3 sf 56003 has 5 sf

48. 2. Don’t count zeros before nonzero numbers. Ex. 0.000078 g has 2 sf 000023.85 has 4 sf Zeros at the end of a number after the decimal are significant. Ex. 83.00 km has 4 sf 0.005520 cm has 4 sf

49. Do not count zeros at the end of a whole number Ex. 2000m has 1 sf

50. Rounding off Numbers Use math rules if number to be considered is 5 or greater, round up if number to be considered is less than 5, round down. Ex. Round this number to 5 sf, then to 3 sf, and finally to 1 sf. 3.515014