Download presentation

Presentation is loading. Please wait.

1
Chemistry an introduction

2
**Why is Chemistry important?**

In our daily lives: New materials New pharmaceuticals New energy sources Food supplies Can you think of anything else?

3
Chemistry Is the science that deals with the materials of the universe and the changes that those materials undergo

4
**Chemical Changes What are some examples of chemical changes?**

Iron rusting Wood burning Food cooking Grape juice fermenting Plants growing How do we know that these are chemical changes?

5
**Steps in the Scientific Method**

Observations Quantitative vs Qualitative Quantitative – measurement involves a number and a unit Formulating Hypotheses Possible explanation for the observation Performing Experiments Gathering new information to decide whether the hypothesis is valid

6
**Quantitative & Qualitative Observations**

Qualitative Quantitative red book 4 quarters round tire 6 wheels wooden desk 24 students metal chair 5 atoms aluminum foil 65°C glass square 2” x 4” x 8” rough board 2 graduated cylinders

7
**Outcomes over the Long Term**

Theory (Model) A set of tested hypotheses that give an overall explanation of some natural phenomenon Natural Law The same observation applies to many different systems Ex. Law of Conservation of Mass

8
Law vs Theory A law summarizes what happens; a theory (model) is an attempt to explain why it happens

9
**The various parts of the Scientific Method**

10
**Problems with the scientific method**

Scientists must be objective when using the scientific method. The scientific method is affected by: Profit motives Religious Beliefs Wars Misinterpretation of Data Budgets Emotions Fads Prejudices Politics Peer Pressure

11
**Scientific Terminology**

What is the difference between a hypothesis and a theory? What is the difference between an observation and a theory? What is the difference between a natural law and a theory?

12
**The Fundamental SI Units**

Physical Quantity Name Abbreviation mass kilogram kg length meter m time second s temperature Kelvin K Electric Current Ampere A Amount of Substance mole mol Luminous Intensity candela cd

13
**SI Prefixes Common to Chemistry**

Unit Abbr. Exponent Mega M 106 Kilo k 103 Deci d 10-1 Centi c 10-2 Milli m 10-3 Micro 10-6 Nano n 10-9 Pico p 10-12

14
**Uncertainty in Measurement**

A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty. Measurements are performed with instruments No instrument can read to an infinite number of decimal places.

15
**Precision and Accuracy**

Accurate and precise Precise, but not accurate Neither accurate not precise Accuracy refers to the agreement between the measure quantity and the accepted value Precision refers to the degree of agreement of several repeated measurements (made in the same manner) to each other.

16
**Types of Error Random Error (Indeterminate Error) –**

Measurement has an equal probability of being high or low Systematic Error (determinate Error) – Occurs in the same direction each time (high or low), often resulting from poor technique or incorrect calibration. This can result in measurements that are precise, but not accurate

17
**Rules for Counting Significant Figures**

Non-zero integers always count as sig. fig. 3456 4 sig figs

18
**Rules for Counting Significant Figures**

Zeros Leading Zeros do not count as sig figs 0.0486 3 sig figs

19
**Rules for Counting Significant Figures**

Zeros Captive Zeros always count as sig figs 16.07 4 sig figs

20
**Rules for Counting Significant Figures**

Zeros Trailing Zeros are significant only if the number contains a decimal point. 9.300 4 sig figs

21
**Rules for Counting Significant Figures**

Exact Numbers have an infinite number of significant figures. 1 inch = 2.54 cm

22
**Practice Counting Significant Figures**

m 17.10 kg 100,890 L 3.29 x103 s cm 3, 200, 000 5 sig figs 4 sig figs 3 sig figs 2 sig figs

23
**Rules for Significant Figures in Mathematical Operations**

Multiplication and Division number of sig figs in the results equals the number of sig figs in the least precise measurement used n the calculation (the one with the lowest number of sig figs). 6.38 x 2.0 = 12.76 13 (2 sig figs)

24
**Practice for Significant Figures in Mathematical Operations**

Answer 23 m2 4.22 g/cm3 0.05 cm2 240 m/s 5870 lb·ft 2.96 g/mL Calculation Calculator Says 3.24 m x 7.0 m 22.68 m2 100.0 g ÷ 23.7 cm3 g/cm3 0.02 cm x cm cm2 710 m ÷ 3.0 s m/s lb x 3.23 ft lb·ft 1.030 g ÷ 2.87 mL g/mL

25
**Rules for Significant Figures in Mathematical Operations**

Addition and Subtraction The number of decimal places in the result equals the number of decimal places in the least precise measurement = 18.7 (3 sig figs)

26
**Practice for Significant Figures in Mathematical Operations**

Calculation Calculator Says 3.24 m m 10.24 m2 100.0 g cm3 76.27 g/cm3 0.02 cm cm 2.391 cm2 713.1 m – s m/s lb lb lb·ft 2.030 mL mL 0.16 g/mL Answer 10.2 m 76.3 g 2.39 cm 709.2 L lb 0.160 mL

Similar presentations

OK

Chemical Foundations. Nature of Measurement Part 1 - number Part 2 - scale (unit) Examples: 20 grams 6.63 x 10 -34 Joule seconds Measurement - quantitative.

Chemical Foundations. Nature of Measurement Part 1 - number Part 2 - scale (unit) Examples: 20 grams 6.63 x 10 -34 Joule seconds Measurement - quantitative.

© 2017 SlidePlayer.com Inc.

All rights reserved.

Ads by Google

Ppt on indian national congress Ppt on primary and secondary data Ppt on industrial development in gujarat ahmedabad Ppt on different types of computer softwares programs Ppt on chapter 3 atoms and molecules class Ppt on polynomials of 90 Ppt on depth first search algorithm examples Ppt on credit policy of a bank Ppt on different types of forests names Slideshare ppt on leadership