Introduction: Matter and Measurement Chapter 1 Introduction: Matter and Measurement Dr. RAJANI SRINIVASAN Tarleton State University

Chemistry the Central Science Chemistry is the central to fundamental understanding of other science and technologies. It shares ties with almost all the branches of science and engineering “Biology, Medical, Agriculture, Pharmacy, Forensic science, Rocket science, Engineering, Environmental Science etc.”

Chemistry ????? Chemistry is the branch of science which studies the properties and behavior of anything that comes in contact with you whether you can see it or not. That anything is “Matter”

Matter We define matter as anything that has mass and takes up space

Matter Matter is composed of the smallest particle called Atom Two or more atoms combine together in a particular shape and angle to form Molecules

Classification of Matter Physical State Composition Element , Compound and mixture Gas, Liquid and solid

States of Matter Gas: has no fixed volume, takes the shape of the container, Can be compressed and expanded Example : water vapor Liquid : Has distinct volume independent of the container , takes the shape o f the container Example : water Solid : Has definite volume and Shape Example : Ice Note: Both Liquid and solid cannot be compressed or expanded.

States of Matter Element : Substances which cannot be decomposed further . They are formed of only one type of atoms. Example : Oxygen , Silicon etc. Currently 118 elements are known Over 90% of Earth’s crust is composed of Oxygen, silicon, aluminium, iron and calcium Oxygen Carbon and hydrogen forms 90% of human body mass

States of matter Compound: Substances formed by the combination of two or more elements . They can be decomposed to their individual components Example : When Oxygen and Hydrogen combine together they form water and when electricity is passed through water they decompose to Oxygen and Hydrogen

States of matter

States of Matter Mixture: Made up of different substances and each substance retains its chemical identity and property. Example : Coffee Mixture Heterogeneous Homogenous Substances which do not have same composition throughout . Eg: rocks and wood Substances which has uniform composition throughout . Eg: Air

Classification of matter

Properties of matter Physical Properties… Chemical Properties… Can be observed without changing a substance into another substance. Boiling point, density, mass, volume, etc. Chemical Properties… Can only be observed when a substance is changed into another substance. Flammability, corrosiveness, reactivity with acid, etc.

Properties of matter Intensive Properties… Extensive Properties… Are independent of the amount of the substance that is present. Density, boiling point, color, etc. Extensive Properties… Depend upon the amount of the substance present. Mass, volume, energy, etc.

Changes in matter Physical Changes These are changes in matter that do not change the composition of a substance. Changes of state, temperature, volume, etc. Chemical Changes (Chemical Reaction) Chemical changes result in new substances. Combustion, oxidation, decomposition, etc.

Chemical Changes In the course of a chemical reaction, the reacting substances are converted to new substances

Separation of mixtures Criteria for separation of mixtures Different methods for different types of mixtures Steps: Determine Use the difference in the properties in the mixture to separate them. Heterogenous Homogenous

Separation of Mixture Example : Heterogeneous mixture of gold and Iron fillings: Properties (Physical):1) Iron magnetic gold non magnetic This property can be used to separate them 2) Chemical Property: Gold dissolves in some acids in which iron does not

Filtration Useful for heterogeneous mixture Solid and liquid mixture can be separated using filtration In filtration, solid substances are separated from liquids and solutions

Distillation Very useful for separating homogenous mixture Used the property of substance to form gasses. Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components Boiling water and salt at 100C water stars evaporating we can condense and get water salt remains in the flask and can be derived from it.

Chromatography This technique separates substances on the basis of differences in solubility in a solvent.

Units of measurement Système International d’Unités There are seven base units from which all the other units are derived A different base unit is used for each quantity In this chapter we will consider Length , mass and temperature

Metric System

Length and Mass SI unit for Length is meter Mass = measure of amount of material in an object. (Note mass is different from weight) [ Weight = Force that is exerted by its on its mass by its gravity] SI unit for mass is Kilo gram

Brain teaser !!!! Question: What is the name of the unit that uses the prefix a) 10 -9 gram – Nano- ng b) 10-6 Second – micro - µg C) 10-3 meter – milimeter- mm

Temperature Temperature is the measure of degree of hotness and coldness of the object Heat always flows from higher temperature to the lower temperature spontaneously. Units for measurement of temperature Celsius Scale - ◦ C Kelvin Fahrenheit

Temperature In scientific measurements, the Celsius and Kelvin scales are most often used. The Celsius scale is based on the properties of water. 0 C is the freezing point of water. 100 C is the boiling point of water

Temperature …… The kelvin is the SI unit of temperature. There are no negative Kelvin temperatures. K = C + 273.15 The Fahrenheit scale is not used in scientific measurements. F = 9/5(C) + 32 C = 5/9(F − 32

Temperature… Note : Zero Kelvin is the lowest attainable temperature = Absolute Zero = -273.15 ◦C Celsius and Kelvin have equal sized units

Derived units Units that are derived from the basic seven SI units are called derived units. Density is a physical property of a substance. It has units (g/mL, for example) that are derived from the units for mass and volume. d = m V

Derived units Speed = Ratio of distance travelled to lapsed time . Speed = meter /second = Length /time

Uncertainty in measurement Exact Numbers – Exact Known value- Dozens Inexact Numbers – Numbers with uncertainty Instrument errors – all the equipment have measuring error Human errors – no two human beings can have same measurement Accuracy – How closely individual measurements agree with the correct value Precision –measure of how closely individual measurements agree with one another

Significant figures When we measure a dime in the balance we report the mass as 2.2405±0.0001g Symbol ± represents the uncertainty in measurements. All the digits in the measured quantity including the uncertain ones are called Significant figures. Greater the significant figures greater certainty implied for the measurement

Significant numbers When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers. All nonzero digits are significant – 234- 3 significant figures Zeroes between two significant figures are themselves significant- 1005- 4 sf; 7.03 = 3 sf Zeroes at the beginning of a number are never significant = 0.45- 2 sf ; 0.03= 1 sf they just indicate the number of decimal Zeroes at the end of a number are significant if a decimal point is written in the number = 2.0 – 2sf

Significant numbers If a number ends with zero but has no decimal then zeros are not considered significant. Note: Exponential notation helps us to understand whether zeros are significant or not Example: 10,300 based on the measurements they can have variable significant figures A) 1.03*10 4 g = 3 sf B) 1.030 * 10 4 g = 4 sf C) 1.0300 * 105g = 5 sf

Significant figures When addition or subtraction is performed, answers are rounded to the least significant decimal place. When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation.

Dimensional analysis Using same units to solve the problems to have accurate answers We use dimensional analysis to convert one quantity to another. Most commonly, dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm

Dimensional analysis Most commonly, dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm version factors to be used 1 in. 2.54 cm or

Dimensional analysis We know 100 cm = 1m or 1m = 1 /100 cm Q: Convert 8.00 meter rod to inches We know 100 cm = 1m or 1m = 1 /100 cm 1inch = 2.54 cm

Problems Use the form of the conversion factor that puts the sought-for unit in the numerator: Given unit   desired unit desired unit given unit Conversion factor

Problems For example, to convert 8.00 m to inches, convert m to cm convert cm to in 8.00 m  100 cm 1 m 1 in. 2.54 cm  315 in.

Conversion involving volume We can also have conversion factor from one measure to different one Example Density D= m/V What will be the mass in grams of 2 cubic inches of gold with density of 19.3 g/cm3 Conversion factor= 19.3g/1cm3 and 1cm3/19.3 g We know 2.54 cm3/1inch3 = 16.39 cm3/1inc3 Mass in grams = 2in3*16.39cm3/1in3* 19.3g/1cm3 = 633 g