One-hour chemistry

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This is chemistry, squeezed into something able to be read in less than 60 minutes (perfect for crammers!). I cannot guarantee that this will cover all of the content for the exam, but it’ll be pretty close. Essentially, I read the book to condense all this information for easy reading. I don’t want to waste words here so here we go:

Chapter 1 - Chemistry: an Introduction

Chemistry is a kind of science that extends to most everything. It specifically deals with substances and their interactions. Or, as the book likes to say, “the science that deals with the materials of the universe and the changes that these materials undergo.” There are implications of chemistry everywhere: in the environment, for medicine, and even in the human body.

Key Concepts:

• Scientific Method:
  • Recognize the problem (observe)
  • Propose possible solutions (hypothesize)
  • Determine which solution is the most reasonable (experiment)
• The five branches of chemistry:
  • Organic chemistry: the chemistry of carbon (and compounds with carbon, i.e. carbohydrates)
  • Inorganic chemistry: the chemistry of compounds without carbon
  • Analytical chemistry: studies the chemical makeup of natural and artifical materials
  • Physical chemistry: the study of the macroscopic and atomic interactions of chemicals in the context of physics. Mathematically intensive and involves thermodynamics, motion, energy, force, time, etc.
  • Biochemistry: the study of chemical processes within living organisms
• Qualitative observation: a description of the appearance of the thing being observed, about one of its qualities (and is not quantifiable, or able to be given a number). An example is saying that the sidewalk appeared to have dirt. You can't say that with a number.
• Quantitative observation: a measurable quantity represented by a number. An example is the temperature at which water boils, being 100℉.
• Theory: Remember TELO (Theory, Explanation, Law, Observation). A theory is defined as a set of thoroughly tested hypotheses that explain an observable behavior in nature. Remember that it is an interpretation of why nature behaves a certain way and not a recordable event like an observation.
• Law: In the LO part of TELO, you can remember that a law is simply a summary of observed behaviors from nature, such as the Law of Conservation of Mass. Laws are the result of recording the same observation for different parts of nature.

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Chapter 2 - Measurements and Calculations

This chapter is about the quantitative measurements we talked about before. Measurements in science often use standard units in a system called SI (International System), but other systems such as the English system and the metric system exist. Remember that the difference between SI and metric is the fact that SI uses kilograms as the base unit for mass, and metric start with grams.

Prefix	Symbol	Value
mega-	M	1,000,000
kilo-	k	1,000
deci-	d	0.1
centi-	c	0.01
milli-	m	0.001
micro-	µ	0.000001
nano-	n	0.000000001

Key Concepts:

• Scientific Notation: uses powers of 10 to describe very large or small numbers. For example:
  • \(1.25\times10^{3} = 1,250\) - In the scientific notation form, the number has 3 significant digits (the power of 10) is not counted.
  • \(9.235\times10^{-5} = 0.00009235\) - Again, the scientific notation form has 4 significant figures.
• Significant figures: There's a problem with doing operations with quantitative measurements of different "precisions" called uncertainty. We correct this by rounding to "significant figures." Here's a reminder of what constitutes significant figures:
  • Any non-zero digits (34, 28, 92 all have 2 sigfigs)
  • Any zeroes in between non-zero digits (101, 302, 0.909 all have 3 sigfigs)
  • Trailing zeroes - ONLY if the number has a decimal (100., .800, 5.30 all have 3 sigfigs)
  Operations:
  • Addition/subtraction: only consider significant digits following the decimal
    • \( 4.02 + 12.483 = 16.50 \) (round to two significant digits after decimal)
    • \( 3.2 + 982 = 985 \) (round to zero significant digits after decimal)
  • Multiplication/division: consider all significant digits
    • \( 600. \times 56.12 = 33700 \) (round to three significant figures)
    • \( 9.3 \times 50 = 500 \) (round to one significant figure)
• Dimensional analysis: In chemistry, since we use many types of units, it is sometimes necessary to convert between these units. We do this by using equivalence statements, or a fraction that equals 1 (multiplicative identity, which does nothing to actual value). An example of an equivalence statement would be \( \frac{12\ pieces}{1 dozen} \). Here are some examples:
  • \( 5\ meters \rightarrow km:\ 5m \times \frac{1 km}{1000 m} = 5000m \)
  • \( 10\ km^{3} \rightarrow m^{3}:\ 10km^{3} \times \frac{1000m}{1km} \times \frac{1000m}{1km} \times \frac{1000m}{1km} = 10000000000m^{3}\)
• Important conversion things to remember:
  • \( 1cm^{3} = 1mL \)
  • \( Temperature_{°Celsius} + 273 = T_{°Kelvin}\)
  • \( Density = \frac{mass}{volume} \)

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Chapter 3 - Matter

Matter, which consists of anything that has mass and takes up space (i.e., pretty much everything), can come as a solid, liquid, or a gas based on its atomic structure. Changes between

Key Concepts:

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Chapter 4 - Chemical Foundations: Elements, Atoms, and Ions

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Chapter 5 - Nomenclature

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Chapter 6 - Chemical Reactions: An Introduction

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Chapter 7 - Reactions in Aqueous Solutions

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Chapter 8 - Chemical Composition

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Chapter 9 - Chemical Quantities