1. Chemistry 2: Scientific Measurement Quantitative measurements give results in a definite form, usually as numbers Qualitative measurements give results in a descriptive nonnumeric form

2. Accuracy and Precision Accuracy is how close a measurement comes to the actual dimension or true value of whatever is measured Precision is concerned with the reproducibility of the measurement

3. Significant Figures in Measurements The significant figures in a measurement include all the digits than can be known precisely plus a last digit that must be estimated. 1. Every nonzero digit is significant. 2. Zeros appearing between nonzero digits are significant. 3. Zeros appearing in front of all nonzero digits are not significant. 4. Zeros at the end of a number and to the right of a decimal point are significant. 5. Use scientific notation whenever possible.

4. The Metric System Lengthmeterm Masskilogramkg Timeseconds Electric currentampereA TemperaturekelvinK Amt of substance molemol Light intensitycandelacd Pressurepascalpa EnergyjouleJ ForcenewtonN CapacitancefaradF

5. Units of Length, meter kilok100010 3 decid1/1010 -1 centic1/10010 -2 millim1/100010 -3 micro10 -6 Nanon10 -9 Picop10 –12

6. Units of Volume, liter A liter is the volume of a cube that is 10 cm on each edge (10cmx10cmx10cm=1000cm 3 =1 L

7. Units of Mass, kilogram A kilogram is the mass of 1 L of water at 4 o C. A gram is defined as the mass of 1 cm 3 of water at 4 o C.

8. Density and Specific Gravity Density is the ratio of the mass of an object to its volume. Density = mass / volume Specific gravity is a comparison of the density of a substance to the density of a reference substance, usually at the same temperature. Specific gravity = density of substance (g/cm 3 ) /density of water (g/cm 3 )

9. Measuring Temperature Temperature is the degree of hotness or coldness of an object. Heat transfer occurs when two objects at different temperatures contact each other. Swedish astronomer Anders Celsius (1701- 1744)—the Celsius scale takes the freezing point of water as 0 o C and boiling point at 100 o C at one atmosphere pressure. Scottish physicist Lord Kelvin (1824-1907)—the Kelvin scale, the freezing point of water is 273 K, and the boiling point is 373 K. The zero point (0 K) on the Kelvin scale is absolute zero. It is –273 o C.

10. Measuring Heat English physicist James Joule (1818- 1889) formulated heat conversion. 1 J = 0.239 cal and 1 cal = 4.18 J 1 calorie is the quantity of heat that raises the temperature of 1 g of pure water 1 o C. 1 kcal = 1000 cal = 1 Cal.

11. Specific Heat Capacity The quantity of heat required to change an object’s temperature by exactly 1 o C is the heat capacity of that object. The specific heat capacity, or simply the specific heat, of a substance is the quantity of heat required to raise the temperature of 1 g of the substance 1 o C.

12. Matter is made up of atoms Joseph Proust, in 1799, observed that water consists of 11% hydrogen and 89% oxygen: in a ratio of 1:8 by mass Antoine Lavoisier, in 1774, discovered the law of conservation of matter John Dalton in 1800s proposed the atomic theory

13. Dalton’s Atomic Theory 1. All matter is made up of atoms 2. Atoms are indestructible and cannot be divided into smaller particles (not true anymore) 3. All atoms of one element are exactly alike, but they are different from atoms of other elements

14. The Methods of Science Repeated observation give rise to hypothesis, which is tested by experiments Repeated experiments either confirm the hypothesis or revise the hypothesis After a hypothesis has been verified my other scientists, it becomes a theory More experiments will give rise to a revised theory When the theory is firmly established without exceptions, it becomes a law

15. The Discovery of Atomic Structure J.J. Thomson, in 1897, discovered the electron using the cathode-ray tube, and that electrons are negatively charged Rutherford’s gold foil experiments showed that atoms have positively charged nuclei Thomson’s atomic model: electrons embedded in a ball of positive charge Nagaoka’s atomic model: electrons orbit around the positive nucleus like planets round the sun

16. Atomic Numbers and Masses Atomic number is the number of protons in the nucleus of an element Mass number is the sum of protons and neutrons in the nucleus. It is the average of all isotopes of an element Isotopes are different forms of an element having the same number of protons but different number of neutrons, hence their masses are different Atomic mass, 1 u = 1/12 the mass of C-12

17. Common Particles of an Atom Particle, symbol, charge, Z, mass in u Protonp + 1 + 11.01 Neutronn 0 011.01 Electrone - 1 - 00.00055

18. Electrons in Motion Niels Bohr proposed that electrons have energies to move round the nucleus Now electrons are considered as particles and waves Electrons move round the nucleus near the speed of light Electrons cannot be accurately located, only the probability of finding it

19. The electromagnetic spectrum All forms of radiant energy can be placed in the electromagnetic spectrum, from radio waves, AM, FM, microwaves, infrared, visible spectrum, UV, X-rays, to gamma rays, which are very energetic and have very short wave-lengths

20. Electron and Light Hydrogen can be shown to consist of many energy levels, which are similar to the rungs on a ladder The electron cloud model of an atom suggests that energy levels are concentric spherical regions of space around the nucleus Electrons in the outermost energy level are called the valence electrons

21. Lewis Dot Diagrams A Lewis dot diagram illustrates valence electrons as dots around the chemical symbol of an element. Each dot represent one valence electron Chemical changes only involve the outermost electrons of an element