Chapter 1 – Introduction to Chemistry What is Chemistry? The Study of matter and energy, and the changes they undergo.Three perspectives 1. Macroscopic: how we observe the world or matter around us.• Physical - Change in form but not in chemical ID- Examples: color, density, odor, and boiling• Chemical - Change in the chemical ID (the ID of the atoms or molecules change)- Examples: flammability or corrosiveness o States of Matter 1. Solid: fixed shape and volume2. Liquid: variable shape and fixed volume3. Gas: variable shape and volume2. Microscopic: All matter is composed of atoms and molecules- Elements are like the building blocks- Atoms are small particles that cannot be made any smaller- Molecules are groups of atoms held together3. Symbolic: what chemist use to represent atoms, molecules, and reactions.- Reactions: A + B → C- Formulas: H2O → water- Elemental symbols: H, O, N, etc. - Ball and stick drawings The Science of Chemistry (observations in Science) Precision vs Accuracy- Precision: how closely individual measurements agree with each other- Accuracy: how closely measurements agree with an “accepted’ value Error - Random errors: fundamental to measurements associated with the limitations of the instrument. Ex) beaker versus a graduated cylinder- Systematic errors: consistently high or low of the accepted value and has an “unknown” bias. Ex) measuring a beaker without being at eye levelModels: largely empirical description- Theory: explanation grounded in a fundamental principle or assumption about the behavior of a system- Law: theories that have been accepted for hundreds of years and are considered self-evident. Numbers and Measurements SI units: international system of units Property UnitMass Kilogram (kg)Time Seconds (s)Distance Meters (m)Electric current Ampere (A)Temperature Kelvin (K)# of particles Moles (mol)Light intensity Candela (cd) Metrix PrefixesGiga (G) 109Mega (M) 106Kilo (k) 103Deci (d) 10-1Centi (c) 10-2Mili (m) 10-3Micro (µ) 10-6Nano (n) 10-9Pico (P) 10-12Femto (f) 10-15 Temperature - The SI unit is Kelvin (Kelvin is on an absolute scale- meaning no negative numbers)- K = ºC + 275.15- ºC = (ºF-32) / 1.8- ºF = (1.8 × ºC) + 32Significant Figures Rules for determining which digits are significant 1) All non-zero digits are significant (5.32 has 3 sig. figs.)2) A zero between 2 significant figures is significant (5.002 has 4 sig. figs.)3) Leading zeros are not significant (0.00532 has 3 sig. figs.) 4) Final zeros in a number with a decimal point are significant (15.0 and 150. both have 3 sig. figs.)5) Final zeros without a decimal point are ambiguous (150 could have 2 OR 3 sig. figs.) Significant figures in Calculations Rule 1 – Multiplication and division - The question is “how many sig figs?” Use the number of sig figs that is least. - 2.99 x 7.3 = [21.827] - The final answer is 22Rule 2 – Addition and subtraction - The question is “how precise?” Carry out the answer to the least precise sig fig.- 29.52 + 3.001 + 7219.5 = [7252.021]- The final answer is 7252.0 Exact numbers: do not limit the number of significant figures1.) Counted objects (20 students, 3 cats)2.) Some conversion factors are exact by definition (2.54 cm = 1