Chapter 1 - Study of Chemistry PDF

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Chapter 1 Chemistry...

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1 Keys to the Study of Chemistry 1.1 Some fundamental definitions Chemistry is the study of matter and its properties, the changes that matter undergoes, and the energy associated with those changes.

The Properties of Matter: The matter is anything that has mass and volume. A substance is a type of matter that has a defined, fixed composition. To identify a substance, chemists observe two types of properties, physical and chemical: -

Physical properties are those that a substance shows by itself, without changing into or interacting with another substance. Some physical properties are color, melting point, electrical conductivity, and density. A physical change occurs when a substance alters its physical form, not its composition. Thus, a physical change results in different physical properties. For example, when ice melts, several physical properties change, such as hardness, density, and ability to flow. But the composition of the sample has not changed: the substance is still water: Water (solid form) → Water (liquid form)

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Chemical properties are those that a substance shows as it changes into or interacts with another substance (or substances). Some examples of chemical properties are flammability, corrosiveness, and reactivity with acids. A chemical change, also called a chemical reaction, occurs when a substance (or substances) is converted into a different substance (or substances): Water --(electric current)--> Hydrogen gas + Oxygen gas.

The Three States of Matter Matter occurs commonly in three physical forms called states: solid, liquid, and gas. -

A solid has a fixed shape that does not conform to the container shape. Particles in solids lie next to each other in a regular, three-dimensional array

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with a definite pattern. A liquid conforms to the container shape but fills the container only to the extent of the liquid’s volume; thus, a liquid forms a surface. Particles in

liquids also lie together but are jumbled and move randomly around one another. -

A gas conforms to the container shape also, but it fills the entire container, and thus, does not form a surface. Particles in the gas usually have great distances between them, as they move randomly throughout the entire container.

Depending on the temperature and pressure of the surroundings, many substances can exist in each of the three physical states, and they can undergo changes in state as well. For example, as the temperature increases, solid water melts to liquid water and then boils to gaseous water (also called water vapor). Similarly, with decreasing temperature, water vapor condenses to liquid water, and the liquid freezes to ice. Thus, a physical change caused by heating can generally be reversed by cooling, and vice versa. This is not generally true for a chemical change. For example, heating iron in moist air causes a chemical reaction that slowly yields the brown, crumbly substance known as rust. Cooling does not reverse this change; rather, another chemical change (or series of them) is required.

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The Central Theme in Chemistry Understanding the properties of a substance and the changes it undergoes leads to the central theme in chemistry: macroscopic properties and behavior, those we can see, are the results of submicroscopic properties and behavior that we cannot see.

The Importance of Energy in the Study of Matter In general, physical and chemical changes are accompanied by energy changes. Energy is often defined as the ability to do work. Essentially, all work involves moving something (W= Force x Displacement): the object doing the work (arm, engine, rock) transfers some of the energy it possesses to the object on which the work is done (book, wheels, ground). The total energy an object possesses is the sum of its potential energy and its kinetic energy. Potential energy is the energy due to the position of the object. Kinetic energy is the energy due to the motion of the object. -

A key concept is that energy is conserved: it may be converted from one form to the other, but it is not destroyed.

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Spontaneous processes (those occurring in a system without the need of continuously providing energy from the surroundings) occur in the direction that is dictated by a decrease of Potential energy of the system

1.2 The Scientific Approach: Developing a Model The scientific method is not a rigid sequence of steps, but rather a dynamic process designed to explain and predict real phenomena.

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1.3 Chemistry Problem Solving Units and Conversion Factors A measured quantity consists of a number and a unit. Conversion factors are used to express a quantity in different units and are constructed as a ratio of equivalent quantities.

Solving Chemistry Problems The problem-solving approach used in this text usually has four parts: 1. Devise a plan for the solution 2. Put the plan into effect in the calculations 3. Check to see if the answer makes sense 4. Practice with similar problems.

1.4 Measurements in Scientific Study Features of SI Units The SI system is based on a set of seven fundamental units, or base units, each of which is identified with a physical quantity. All other units, called derived units, are combinations of these seven base units.

  For quantities that are much smaller or much larger than the base unit, we use decimal prefixes and exponential (scientific) notation.

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SI Units in Chemistry -

Length: The SI base unit is the meter - m (derived by light’s velocity); biological cells are often measured in micrometers (1µm = 10-6 m); on the atomic-size scale, nanometers and picometers are used (1nm= 10-9  m; 1 pm= -12 10 m); atomic diameters are around 200 pm. An older unit still in use is the

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angstrom (1 Å= 10-10m). Volume: Any sample of matter has a certain volume, the amount of space that the sample occupies. The SI unit of volume is the cubic meter (m3 ). In chemistry, the most important volume units are non-SI units, the liter (L) and the milliliter (mL) - note the uppercase L- ; physicians and other medical practitioners measure body fluids in cubic decimeters (dm3 ), which is equivalent to liters: 1L= 1dm3= 10-3m3 . As the prefix milli- indicates, 1 mL is

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1/1000 of a liter, and it is equal to exactly 1 cubic centimeter (cm3). Mass: The mass of an object refers to the quantity of matter it contains. Other base units are the gram (10-3  kg) and the milligram (10-3  g). Remember that the terms mass and weight have distinct meanings: its mass is constant while its weight depends both on its mass and the strength of the local gravitational

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field pulling on it. Density: The density (d) of an object is its mass/volume. Because volume may change with temperature, density may change also. The SI unit of density is the kilogram per cubic meter (kg/m3 ), but in chemistry, density is typically given in units of g/L (g/dm3) or g/mL (g/cm3)

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Temperature: Temperature (T) is a measure of how hot or cold a substance is relative to another substance (Heat is the energy that flows between objects that are at different temperatures, which is very different). It is related to the direction of that energy flow: when two objects at different temperatures touch, energy flows from the one with the higher temperature to the one with the lower temperature until their temperatures are equal. In the laboratory, the most common means for measuring temperature is the thermometer, a device that contains a fluid that expands when it is heated. The three temperature scales most important for us to consider are the Celsius (°C), the Kelvin (K), and the Fahrenheit (°F) scales. The SI base unit of temperature is the kelvin (K), also known as the absolute scale, and it is preferred in all scientific work, although the Celsius scale is used frequently. In the United States, the Fahrenheit scale is still used for weather reporting, body temperature, and other everyday purposes. The three scales differ in the size of the unit and/or the temperature of the zero point. The zero point in the Kelvin scale, 0 K, is called absolute zero and equals 273.15°C and, in the Kelvin scale, all temperatures have positive values. Conversion formulas: T (in K) = T (in °C) + 273.15 T (in °F) = [9/5 T (in °C)] + 32

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Time: The SI base unit of time is the second (s). The standard second is defined by the number of oscillations of microwave radiation absorbed by cooled gaseous cesium atoms in an atomic clock; precisely 9,192,631,770 of these oscillations are absorbed in 1 second. In the laboratory, we study the speed (or rate) of a reaction by measuring the time it takes a fixed amount of substance to undergo a chemical change. The range of reaction rates is enormous: a fast reaction may be over in less than a nanosecond (109 s), whereas slow ones, such as rusting or aging, take years. Chemists now use lasers to study changes that occur in a few picoseconds (1012 s) or femtoseconds (1015 s).

Extensive and Intensive properties

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1.5 Uncertainty in Measurement: Significant We can never measure a quantity exactly, because measuring devices are made to limited specifications and we use our imperfect senses and skills to read them: every measurement includes some uncertainty.

Determining Significant Figures When you take measurements or use them in calculations, you must know the number of digits that are significant. In general, all digits are significant, except zeros that are not measured but are used only to position the decimal point. Here is a simple procedure that applies this general point: 1. Make sure that the measured quantity has a decimal point. 2. Start at the left, and move right to the first nonzero digit. 3. Count that digit and every digit to its right as significant. A complication may arise with zeros that end a number. Zeros that end a number and lie either after or before the decimal point are significant; thus, 1.030 mL has four significant figures, and 5300. L has four significant figures also. If there is no decimal point, as in 5300 L, we assume that the zeros are not significant; exponential notation is needed to show which of the zeros, if any, were measured and therefore are significant. A terminal decimal point is used to clarify the number

of significant figures; thus, 500 mL has one significant figure, but 5.00*102 mL, 500. mL, and 0.500 L have three.

Significant Figures in Calculations Measurements often contain differing numbers of significant figures. In a calculation, we keep track of the number of significant figures in each quantity so that we don’t claim more significant figures (more certainty) in the answer than in the original data. If we have too many significant figures, we round off the answer to obtain the proper number of them. The general rule for rounding is that the least certain measurement sets the limit on certainty for the entire calculation and determines the number of significant figures in the final answer. Suppose you want to find the density of a new ceramic. You measure the mass of a piece on a precise laboratory balance and obtain 3.8056 g; you measure its volume as 2.5 mL by displacement of water in a graduated cylinder. The mass has five significant figures, but the volume has only two, therefore, you report the answer with two significant figures 1.5 g/mL.  Rules for rounding off arithmetic operations: 1. For multiplication and division. The answer contains the same number of significant figures as in the measurement with the fewest significant figures. ex:

(two significant figures).

2. For addition and subtraction. The answer has the same number of decimal places as there are in the measurement with the fewest decimal places. ex:

(one decimal place).

 Rules for Rounding Off 1. If the digit removed is more than 5, the preceding number is increased by 1: 5.379 rounds to 5.38 if three significant figures are retained and to 5.4 if two significant figures are retained. 2. If the digit removed is less than 5, the preceding number is unchanged: 0.2413 rounds to 0.241 if three significant figures are retained and to 0.24 if two significant figures are retained. 3. If the digit removed is 5, the preceding number is increased by 1 if it is odd and remains unchanged if it is even: 17.75 rounds to 17.8, but 17.65 rounds to 17.6. 4. Always carry one or two additional significant figures through a multistep calculation and round off the final answer

 Exact numbers Some numbers are called exact numbers because they have no uncertainty associated with them. Some exact numbers are part of a unit definition: there are 60 minutes in 1 hour, 1000 micrograms in 1 milligram, and 2.54 centimeters in 1 inch.

Precision and Accuracy Precision and accuracy are two aspects of certainty. Precision refers to how close the measurements in a series are to one another. Accuracy refers to how close a measurement is to the actual value.Precision and accuracy are linked with two common types of error: 1. Systematic error produces values that are either all higher or all lower than the actual value. Such error is part of the experimental system, often caused by a faulty measuring device or by a consistent mistake in taking a reading (accurate measurements have low systematic error). 2. Random error, in the absence of systematic error, produces values that are higher and lower than the actual value. Random error always occurs, but its size depends on the measurer’s skill and the instrument’s precision (precise measurements have low random error).

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