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**Measurement in Chemistry (and elsewhere)**

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**Types of observations Qualitative**

Properties that can be observed and described that do not involve measurement. If they do refer to quantities, they are vague (ie fast, hot, large etc…) Quantitative Properties that can be observed and described numerically and which result from measurement.

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**Commonly measured values in chemistry**

Mass (grams) Volume (liters) Length (meters) Temperature (degrees Celsius or degrees Kelvin) Time (seconds) Pressure (atmospheres) Concentration (percent, molar)

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**Length, Mass and Volume Length - distance between two points**

Mass - amount of matter in an object (weight is dependent upon the force of gravity on the object) Volume - the amount of space an object occupies

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Length, Mass and Volume Fig 2.2

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**Volume is derived from length**

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**Some units of measurement**

Metric Base Unit SI Unit Length meter (m) meter Mass gram (g) kilogram (kg) Volume liter (L) meter3 Temperature Celsius (C) Kelvin (K)

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**Common metric prefixes**

Table 2.1

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**Common metric prefixes**

1000 base / 1 kilo g / 1 kilogram x 103 100 centi / 1 base 100 centimeters/ 1 meter 1 x 102 1000 milli / 1 base millimeters/ 1 meter 1 x 103 1,000,000 micro / 1 base 1,000,000 micrometers/1 meter 1 x 106

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**Converting between units (Dimensional analysis – factor label method)**

Given unit x (Desired unit) = Desired unit (Given unit) 12.4 kg x (1000 g) = 12,400 g (1kg) 1265 mm x (1 m) = m (1 x 103 mm)

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**Exact and inexact numbers**

Exact numbers No uncertainty to their value Value is known exactly Defined Conversions within a systems Inexact numbers Uncertainty of their true value Measured Conversions between different systems

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**Expressing numbers in scientific notation**

Why do it? How to enter them into your calculator 1.5 x 1023 EXP (or EE) 2.67 x 10-16 EXP (or EE) /

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Making measurements Accuracy: How close a measured value is to the true value Precision: How close multiple measured values are to each other

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**There is estimation (and therefore uncertainty) in all measurements**

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Significant figures The digits in a measurement that are known with certainty, plus the single estimated digit Only applies to measured (inexact) values Does not apply to defined (exact) values

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**Measured Values What figures (digits) are significant?**

(not applied to defined or exact values such as conversions within the same system) Non zeros are significant Zeros between non zeros are significant Zeros at the beginning are not significant Zeros at end after decimal are significant Zeros at end before the decimal depend Three ways to represent these zeros

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**How many significant figures are in these measured values?**

cm L mm 100 kg x mg

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**Rules for working with measured values**

Since there is uncertainty in measurement, we risk “amplifying” the uncertainty when we add, subtract, multiply and divide measured values So…. There are rules for working with measured values

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**Calculations involving measured values**

Multiplying and dividing: Answer can have no more total sig. figs. than the starting value with the fewest total sig. figs. Adding and Subtracting: Answer can have no more sig. figs. after the decimal than any original number

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**Dimensional analysis helps solve conversion problems**

What are you starting with? What do you need to convert it into? What conversion factor(s) do you need? Must know conversions within the metric system. Must know other conversions we will identify. Do not have to memorize conversions between systems.

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**English/Metric conversions (Table 2.2)**

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**Density Mass of material per given volume Commonly: grams/mL SI: kg/m3**

Density is a conversion factor for converting between mass and volume grams (mL/g) = milliliter milliliter (g/mL) = grams

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Temperature scales K = C C = K

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**Calories and specific heat**

calorie: amount of heat 1 cal raises 1 g of water 1° C 60 Calories = 60 kcal = 60,000 calories Specific heat of any substance Amount of heat (in calories) required to raise 1 gram of the substance 1° C

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