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Chapter 1 Basic Concepts of Chemistry

Chapter 1 Basic Concepts of Chemistry. Classifying Matter: States of Matter. Molecules in a Solid, Liquid, & Gas. Classifying Matter. Classifying Matter. Matter and its Representation. What we observe…. To what we can ’ t see!. Chemical symbols allow us to connect…. Elements.

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Chapter 1 Basic Concepts of Chemistry

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  1. Chapter 1Basic Concepts of Chemistry

  2. Classifying Matter: States of Matter

  3. Molecules in a Solid, Liquid, & Gas

  4. Classifying Matter

  5. Classifying Matter

  6. Matter and its Representation What we observe… To what we can’t see! Chemical symbols allow us to connect…

  7. Elements • The elements are recorded on thePERIODIC TABLE • There are 117 recorded elements at this time. • The Periodic table will be discussed further in chapter 2.

  8. Chemical compounds are composed of two or more atoms. Chemical Compounds

  9. Chemical Compounds Ionic Compound Iron pyrite (FeS2) Molecule: Ammonia (NH3)

  10. Chemical Compounds • All Compounds are made up of molecules or ions. • A molecule is the is the smallest unit of a compound that retains its chemical characteristics. • Ionic compounds are described by a “formula unit”. • Molecules are described by a “molecular formula”.

  11. Molecular Formula • A moleculeis the smallest unit of a compound that retains the chemical characteristics of the compound. • Composition of molecules is given by a molecular formula. C8H10N4O2 - caffeine H2O

  12. Physical Properties • Some physical properties: • Color • State (s, g or liq) • Melting and Boiling point • Density (mass/unit volume) • Extensive properties(mass) depend upon the amount of substance. • Intensive properties(density) do not.

  13. The Nature of Matter Chemists are interested in the nature of matter and how this is related to its atoms and molecules. Gold Mercury

  14. Crystal Structures: Ag & Au

  15. Physical Properties O H H Physical properties are a function of intermolecular forces. Water (18 g/mol) liquid at 25oC Methane (16 g/mol) gas at 25oC H • Water molecules are attracted to one another by • “hydrogen bonds”. • Methane molecules only exhibit week “London • Forces”. C H H H

  16. Physical Properties Physical properties are affected by temperature (molecular motion). The density of water is seen to change with temperature.

  17. Physical Properties Mixtures may be separated by physical properties:

  18. Chemical Properties • Chemical properties are really chemical changes. • The chemical properties of elements and compounds are related to periodic trends and molecular structure.

  19. Hyrdogen Balloon Ignited

  20. Chemical Properties A chemical property indicates whether and sometimes how readily a material undergoes a chemical change with another material. For example, a chemical property of hydrogen gas is that it reacts vigorously with oxygen gas.

  21. Reaction of Al + Br2

  22. A Chemist’s View of Water Macroscopic H2O (gas, liquid, solid) Symbolic Particulate 2 H2(g) + O2 (g)  2 H2O(g)

  23. Water Droplet/Molecules

  24. Energy: Some Basic Principles Energy can be classified as Kinetic or Potential. • Kinetic energyis energy associated with motion such as: • The motion at the particulate level (thermal energy). • The motion of macroscopic objects like a thrown baseball, falling water. • The movement of electrons in a conductor (electrical energy). • Wave motion, transverse (water) and compression (acoustic). Matter consists of atoms and molecules in motion.

  25. Energy: Some Basic Principles Potential energyresults from an object’s position: • Gravitational: An object held at a height, waterfalls. • Energy stored in an extended spring. • Energy stored in molecules (chemical energy, food) • The energy associated with charged or partially charged particles (electrostatic energy) • Nuclear energy (fission, fusion).

  26. Units of Measure Science predominantly uses the “SI” (System International) system of units, more commonly known as the “Metric System”.

  27. Units of Measure The base units are modified by a series of prefixes which you will need to memorize.

  28. Temperature Units Temperature is measured in the Celsiusan the Kelvintemperature scale.

  29. Length, Volume, and Mass Unit conversions: How many picometers are there in 25.4 nm? How many yards?

  30. Length, Volume, and Mass The base unit of volume in the metric system is the liter. 1 L = 103 mL 1 mL=1 cm3 1 cm3 = 1 mL

  31. Length, Volume, and Mass The base unit of mass in the metric system is the gram. 1kg = 103g

  32. Energy Units Energy is confined as the capacity to do work. The SI unite for energy is the joule (J). Energy is also measured in calories (cal) 1 cal = 4.184J A kcal (kilocalorie) is often written as Cal. 1 Cal =103 cal

  33. Mathematics of Chemistry Exponential or Scientific Notation: Most often in science, numbers are expressed in a format the conveys the order of magnitude. 3285 ft = 3.285  103 ft 0.00215kg = 2.15  103 kg

  34. Exponential or Scientific Notation Exponential part 1.23  104 Coefficient or Mantissa (this number is 1 and <10 in scientific notation Base Exponent

  35. Mathematics of Chemistry Significant figures: The number of digits represented in a number conveys the precision of the number or measurement. A mass measured to  0.1g is far less precise than a mass measured to  0.0001g. 1.1g vs. 1.0001g (2 sig. figs. vs. 5 sig. figs) In order to be successful in this course, you will need to master the identification and use of significant figures in measurements and calculations!

  36. Counting Significant Figures • All non zero numbers are significant • All zeros between non zero numbers are significant • Leading zeros are NEVER significant. (Leading zeros are the zeros to the left of your first non zero number) • Trailing zeros are significant ONLY if a decimal point is part of the number. (Trailing zeros are the zeros to the right of your last non zero number).

  37. Determining Significant Figures zeros written explicitly behind the decimal are significant… Determine the number of Sig. Figs. in the following numbers 1256 4 sf 1056007 7 sf 0.000345 3 sf 0.00046909 5 sf 4 sf 770.0 0.08040 4 sf

  38. Sig. Figures in Calculations Multiplication/Division The number of significant figures in the answer is limited by the factor with the smallest numberof significant figures. Addition/Subtraction The number of significant figures in the answer is limited by the least precise number(the number with its last digit at the highest place value). NOTE: counted numbers like 10 dimes never limit calculations.

  39. Sig. Figures in Calculations Determine the correct number of sig. figs. in the following calculation, express the answer in scientific notation. from the calculator: 4 sf 2 sf 4 sf = 1996.501749 10 sf 23.50  0.2001  17 Your calculator knows nothing of sig. figs. !!!

  40. Sig. Figures in Calculations Determine the correct number of sig. figs. in the following calculation, express the answer in scientific notation. 1.996501749  103 in sci. notation: 2.0  103 Rounding to 2 sf:

  41. Sig. Figures in Calculations Determine the correct number of sig. figs. in the following calculation:  12.6 391 + 156.1456

  42. Sig. Figures in Calculations no digits here To determine the correct decimal to round to, align the numbers at the decimal place: One must round the calculation to the least significant decimal.  12.6 391 +156.1456 391 12.6 +156.1456

  43. Sig. Figures in Calculations one must round to here 391 -12.6 +156.1456 round to here (units place) 534.5456 (answer from calculator) Answer: 535

  44. Problem Solving and Chemical Arithmetic Dimensional Analysis: Dimensional analysis converts one unit to another by using conversion factors (CF’s). The resulting quantity is equivalent to the original quantity, it differs only by the units. unit = new unit  conversion factor

  45. Problem Solving and Chemical Arithmetic Dimensional Analysis: Dimensional analysis converts one unit to another by using conversion factors (CF’s). Conversion factors come from equalities: 1 m = 100 cm 1 m 100 cm or 100 cm 1 m

  46. Examples of Conversion Factors Exact Conversion Factors:Those in the same system of units 1 m = 100 cm Use of exact CF’s will not affect the significant figures in a calculation.

  47. Examples of Conversion Factors Inexact Conversion Factors:CF’s that relate quantities in different systems of units 1.000 kg = 2.205 lb SI units British Std. (4 sig. figs.) Use of inexact CF’s will affect significant figures.

  48. Problem Solving and Chemical Arithmetic Most importantly, before you start... PUT YOUR CALCULATOR DOWN! Your calculator wont help you until you are ready to solve the problem based on your strategy map.

  49. Problem Solving and Chemical Arithmetic Example: How many meters are there in 125 miles? First: Outline of the conversion:

  50. Problem Solving and Chemical Arithmetic Example: How many meters are there in 125 miles? First: Outline of the conversion: m miles  ft  in  cm  Each arrow indicates the use of a conversion factor.

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