General Chemistry I CHEM-101

Introduction Chemistry: Branch of Science which deals with study of composition of matter, properties of matter, changes in matter and law and principles under which these changes occur. Or Chemistry is the science of atoms, their structures, their combination and their interactions. Science: A process for understanding nature and its changes to explain phenomena of the physical world. Matter: Anything which as some mass and occupy some space. Examples: water, Gold, NaCl, sugar, Air etc.

1.1. Types of matter

Colloid - Mixture in which the different parts can be seen, but the particles do not separate and settle out. Some colloids particles are sometimes small enough they cannot be seen to the unaided eye. These particles can be detected however by scattering light due to the Tyndall effect. Suspension - Mixture in which particles of one substance settle to the bottom. ( Water and Oil)

Homogeneous mixture: Having visibly indistinguishable parts Homogeneous mixture: Having visibly indistinguishable parts. Physical properties are the same throughout the material. A homogeneous mixture a solution (example: vinegar). Heterogeneous mixtures: Having visibly distinguishable parts. Physical properties are different at different points in a material (example: bottle of ranch dressing). 6

Definitions Pure substance: is one with constant composition. Pure substances can be isolated by separation techniques – distillation, filtration, chromatography. Compound: is a substance with constant composition that can be broken down into elements by chemical processes. Example: electrolysis of water produces hydrogen and oxygen. Physical change: is a change in the form of a substance but not in its chemical composition. E.g. conversion of ice into water. Chemical change: when a given substance becomes a new substance or substances with different properties and different composition. Burning of wood or gas. 7

Compound

States of Matter

There are four types of crystalline solids - Ionic solids-- These substances have a definite melting point and contain ionic bonds. An example would be sodium chloride (NaCl). Covalent solids -- These substance appear as a single giant molecule made up of an almost endless number of covalent bonds. An example would be graphite. Molecular solids are represented as repeating units made up of molecules. An example would be ice. Metallic solids are repeating units made up of metal atoms. The valence electrons in metals are able to jump from atom to atom.

1.2. Measurement; quantities and units Making measurements is essential for all sciences. In chemistry, we generally use measurements of mass, temperature, time, amount of substance (number of moles),… If measurements were to be useful, a standard system of units has to be adopted. There are 2 majors standard systems of units which are adopted around the world:

non-standard units (past)

Seer (unit) Tola (unit)

(Le System International in French) or the SI system. Units of Measurement The metric system: used by most of the world The English system: used in the USA In 1960, an international agreement set up a system of units called the International System (Le System International in French) or the SI system. The SI system is based on the metric system and units are from the metric system.

Note that the volume V, which is a very important physical quantity in chemistry, is not a fundamental unit but is derived from length. A cube that measures 1 m on each edge has a volume of 1 m3.

Precision and Accuracy Precision: Determines how closely measurements agree with each other, Reflects the reproducibility of a given type of measurement. 1st series of measurements: 34, 35, 37, 37, 38 2nd series of measurements: 30, 35, 40, 42, 47 The precision of the 1st series is better than the 2nd series. Accuracy: Determines the agreement of a particular value with the true value (standard value). Example: True Value = 36.0 Average = Sum of all the values / number of values example: 34, 35, 37, 37, 38 sum of all the values = 181 number of values = 5 average = 36.2 true value = 36.0  good accuracy. 22

Precision and Accuracy No Precision No accuracy Precision but not accuracy Accuracy The Difference between Precision and Accuracy 23

Errors Random Error (indeterminate error): A measurement has an equal probability of being high or low. This type of error occurs in estimating the value of the last digit of measurement. Systematic Error (Determinate error): This type of error occurs in the same direction each time. It is either always high or always low, often resulting from poor technique. 26

1.3. Uncertainties in measurements; significant figures All non-zero digits are significant figures: 1234  4 significant figures Zeros between non-zero digits (captive zeros) are significant figures: 205  3 significant figures Zeros beyond decimal point at the end of the number (trailing) are significant figures: 0.24000  5 significant figures Zeros preceding the first non-zero digit in a number are not significant figures: 0.00453  3 significant figures Zeros at the end of whole numbers are not significant figures unless you are given information to the contrary: 3400  2 significant figures, 34.00 X 102  4 significant figures, 3400.  4 significant figures. 27

How many significant figures are in each of the following? 12  2 significant figures (S.F.) 1098  4 S.F. 2001  4 S.F. 2.001 x 103  4 S.F. 0.0000101  3 S.F. 1.01 x 10-5  3 S.F. 1000.  4 S.F. (because of the decimal point). 22.04030  7 S.F. 28

Rules for Significant Figures in Mathematical Operations Multiplication and division: The number of significant figures in the result is the same as that in the quantity with the smallest number of significant figures. Ex.: 4.56 x 1.4 = 6.38  6.4 (3 S.F.) (2 S.F.) (3 S.F.) (2 S.F.) The product should have only two significant figures since 1.4 has two significant figures. Addition and subtraction: The result has the same number of decimal places as the least precise measurement used in the calculation. Ex.: 12.11 + 18.0 + 1.013  31.128  31.1 The correct result is 31.1, since 18.0 has only one decimal place. 29

Rules for Rounding of Data In a series of calculations, carry the extra digits through to the final result, then round. If the digit to be removed; a. Is less than 5, the preceding digit stays the same. Example. 1.33 rounds to 1.3. b. Is equal to or greater than 5, the preceding digit is increased by 1. Example. 1.36 rounds to 1.4. 31