Download presentation

1
**Unit 1-Chemistry and Measurement**

Chemistry studies matter and the changes matter undergoes. Matter is anything that occupies space and has mass. Sugar Water Gold Dura GOB

2
**All Matter Classifying Matter Pure Substances Mixtures Elements**

NO YES Can it be separated by a physical process? Pure Substances Mixtures Can it be broken down into simpler ones by chemical means? NO YES Elements Compounds 1.6

3
**Heterogeneous mixture – composition is not uniform throughout.**

A mixture is a combination of two or more substances in which the substances retain their distinct identities. Homogenous mixture – composition of the mixture is the same throughout. soft drink, milk, solder Heterogeneous mixture – composition is not uniform throughout. cement, iron filings in sand Dura GOB

4
**The Three Physical States of Matter**

solid liquid gas Dura GOB

5
**Change: Physical or Chemical?**

A physical change does not alter the composition or identity of a substance. ice melting sugar dissolving in water A chemical change alters the composition or identity of the substance(s) involved. hydrogen burns in air to form water Dura GOB

6
**Extensive and Intensive Properties**

An extensive property of a material depends upon how much matter is is being considered. mass length volume An intensive property of a material does not depend upon how much matter is being considered. density temperature color Dura GOB

7
**International System of Units (SI)**

Dura GOB

8
Dura GOB

9
**Uncertainty in Measurement: Significant Figures**

Every measurement includes some uncertainty. Measuring devices are made to limited specifications and we use our imperfect senses and skills. Significant figures: The digits we record in a measurement that include both certain digits and the first uncertain or estimated one. The number of significant figures in a measurement depends on the design or specifications of measuring device. Dura GOB

10
**Aspects of Certainty: precision and accuracy**

Accuracy – how close a measurement is to the true value. Precision – how close a set of measurements are to each other accurate & precise precise but not accurate not accurate & not precise Dura GOB

11
**Precision and Accuracy are linked to two common types of error:**

Random Errors – always occur. Random errors result in some values that are higher and some values that are lower than the actual value. These errors depend on the measurer’s skill and the instruments readability. Precise measurements have low random error. Systematic Errors occur due to poor experimental design or procedure or faulty equipment. Accurate measurements have low systematic error and low random error. Dura GOB

12
**Significant Figures in Measurements**

Any digit that is not zero is significant 1.234 kg significant figures Zeros between nonzero digits are significant 606 m significant figures Zeros to the left of the first nonzero digit are not significant 0.08 L significant figure If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg significant figures If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant g 4 significant figures Dura GOB

13
**How many significant figures are in each of the following measurements?**

24 mL 2 significant figures 3001 g 4 significant figures m3 3 significant figures 6.4 x 104 molecules 2 significant figures 560 kg 2 significant figures Dura GOB

14
**Significant Figures in Calculations**

Addition or Subtraction The answer cannot have more digits to the right of the decimal point than any of the original numbers. 89.332 1.1 + 90.432 one significant figure after decimal point round off to 90.4 3.70 0.7867 two significant figures after decimal point round off to 0.79 Dura GOB

15
**Significant Figures in Calculations**

Multiplication or Division The number of significant figures in the result is set by the original number that has the smallest number of significant figures 4.51 x = = 16.5 3 sig figs round to 3 sig figs 6.8 ÷ = = 0.061 2 sig figs round to 2 sig figs Dura GOB

16
**Significant Figures Exact Numbers**

Numbers from definitions or numbers of objects are considered to have an infinite number of significant figures The average of three measured lengths; 6.64, 6.68 and 6.70? 3 = = 6.67 = 7 Because 3 is an exact number Dura GOB

17
**Scientific Notation The number of atoms in 12 g of carbon:**

602,200,000,000,000,000,000,000 6.022 x 1023 The mass of a single carbon atom in grams: 1.99 x 10-23 N x 10n N is a number between 1 and 10 n is a positive or negative integer Dura GOB

18
**Scientific Notation is used to avoid ambiguity in the number of significant figures in calculations.**

Multiplication (4.0 x 10-5) x (7.0 x 103) = (4.0 x 7.0) x (10-5+3) = 28 x 10-2 = 2.8 x 10-1 Multiply N1 and N2 Add exponents n1 and n2 Division 8.5 x 104 ÷ 5.0 x 109 = (8.5 ÷ 5.0) x = 1.7 x 10-5 Divide N1 and N2 Subtract exponents n1 and n2 Dura GOB

19
**Scientific Notation Exercises**

move decimal left move decimal right n > 0 n < 0 = x 102 = 7.72 x 10-6 Addition or Subtraction Write each quantity with the same exponent n Combine N1 and N2 The exponent, n, remains the same 4.31 x x 103 = 4.31 x x 104 = 4.70 x 104 Dura GOB

20
**Density 1 g/cm3 = 1 g/mL = 1000 kg/m3 density = mass volume d = m V**

A piece of platinum metal with a density of g/cm3 has a volume of 4.49 cm3. What is its mass? How many significant figures should the final answer have? m = d x V = g/cm3 x 4.49 cm3 = 96.5 g Dura GOB

22
**Error and Uncertainty Analysis**

Percent error = I theoretical – experimental I x 100% Theoretical Note: Numerator is an absolute value so that percent error is always positive Dura GOB

23
Uncertainty Analysis Uncertainty should be reported to only one significant digit Proper Improper Fractional uncertainty = uncertainty measurement Dura GOB

24
**Fractional uncertainties are often converted to percentage uncertainties by multiplying them by 100**

Rule 1: When adding or subtracting measurements the uncertainty of the result is the sum of the terms used [ cm] + [ cm] = cm Rule 2: When multiplying or dividing measurements, the fractional uncertainty in the result is the sum of the fractional uncertainties in the factors used Dura GOB

25
Example: [ cm] x [ cm] Fractional uncertainties: 0.1 / 41.7 = , 0.1 / 12.1 = Product: 41.7 cm x 12.1 cm = 505 cm2 Uncertainty: = x 505cm2 = 5cm2 Answer: cm2 Dura GOB

26
Dura GOB

Similar presentations

© 2021 SlidePlayer.com Inc.

All rights reserved.

To make this website work, we log user data and share it with processors. To use this website, you must agree to our Privacy Policy, including cookie policy.

Ads by Google