Chem 101 1st Edition Lecture 4 Outline of Last Lecture I. Introduction to Significant Figures, Scientific Notation, Accuracy and Precision. Outline of Current Lecture II. The Law of Conservation of MassIII. Dalton’s Atomic TheoryIV. The Law of Definite ProportionsV. The Law of Multiple ProportionsVI. Atomsa. IsotopesVII. ElectronsCurrent LectureI. The Law of Conservation of Mass: States that mass is neither created nor destroyed in a chemical reaction but merely changes form. a. The additive mass of all reactants = total massb. Consider: CaO + CO ₂  CaCO ₃You take the mass of Ca: 40.08 C: 12.01O : 16.00 O : (2x) 16.00 = 32₂ 56.08 amu + 44.01 amu Gives you a total mass of: 100.09 amuc. AMU: Atomic Mass UnitsII. Dalton’s Atomic Theory:a. All matter is made up of tiny, indivisible particles called atomsb. Atoms cannot be created, destroyed, or transformed into other atoms in a chemical reactionc. All atoms of a given element are identicald. Atoms combine in simple, whole-number ratios to form compoundsIII. The Law of Definite Proportions: States that a compound is always made up of the samerelative masses of the elements that compose it.a. Example: 2 atoms of H per atom of O : will always form H₂O (water)IV. The Law of Multiple Proportions: states that any time two or more elements combine indifferent ratios, different compounds are formed.These notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.a. Example: Laughing gas, which has a chemical formula of N₂O is quite different from the gas NO₂, a byproduct of explosions and burning diesel fuel. V. Atoms: electrically neutral, spherical entities, composed of a positively charged central nucleus surrounded by one or more negatively charged electrons. a. Atomic Nucleolus: Consists of protons (+) and neutrons (0)i. Protons: sometimes written as [ p ]⁺ has a mass of: 1.0073 amuii. Neutrons: neutrally charged [ n ] has a mass of : ᵒ 1.0087 amuiii. Electron: Sometimes written as [ e¯ ] has a mass of: .0005486 amu1. more than a thousand times smaller than [ p⁺ ] or [ nᵒ ]b. Atoms vary in size depending on electron repulsion.c. Each element is defined by how many protons it has in its nucleus. The protons give the element its identity. i. Z : stands for the Atomic Number/ the number of p⁺ in an elementii. The number of electrons = the number of protons in neutral elements.1. Ex: Li : Z= 3 meaning Li has 3 protons and 3 electrons2. If an element doesn’t have a + or – symbol associated with it it’s a neutral element. d. Isotopes: species with the same # of protons but different numbers of neutrons.i. Example: Hydrogen Isotopesii. Let us consider Cl ( Chlorine ) isotopes: 1. Cl 35 makes up 75.78% of all Cl found in naturea. Its atomic mass is 34.969 amu2. Cl 37 makes up 24.22% of all Cl found in naturea. Its atomic mass is 36.966 amu The atomic mass seen on the periodic table is a calculated average of all Cl and Cl isotopes that occur in nature. To calculate the atomic mass of any element:1. You first determine how many isotopes are associated with theelement you wish to calculate2. Take the atomic mass of one isotope and multiply it by the % of abundance in nature3. You do this for all the existing isotopes of that element4. Once you have the sums of the mass of the isotopes and the %of abundance you add them all up and divide by the number of isotopes. This is a chart of some of the most abundant isotopes obtained from www.sepscience.com as an example for your reference. You do not have to memorize this chart. VI. The Electron: The electrons are “where the chemistry happen”. All chemical reactions involve electronsa. Electrons are arranged in Orbitals or different levels of energy.i. Let us consider a Hydrogen. Hydrogen has one electron that generally inhabits the first energy level. This is considered to be a ground state hydrogen atom.ii. If you excite the electron and it absorbs energy the electron may move to the 2nd or 3rd energy levels. If this happens it is generally brief and unstable. The electron will release its energy and return to the 1st energy level. We call this jumping to higher energy levels an excited state Hydrogen atom. b. Let us now consider other elements. Because electrons are all negatively chargedthey repel themselves. However they are all attracted to the positively charged protons in the nucleus. Because every other element has more than one electronwe will use a different notation: Energy sublevelsc. Orbitals: are 3 dimensional regions around the nucleus where the electron has the highest probability of being located. Because the next couple concepts are visually based I will be using charts and pictures to better convey the material.d. Orbitals come in many complex shapes and configurations. You will need to memorize the s and p orbitals shown above.Rules for Electron Orbitals: 1. A single orbital can contain a maximum of 2 electrons2. When filling electrons in the p, d, and f subsets, each orbital gets a single electron before any orbital of the subset receives a second electron. e. The s orbital is always a sphere and stands for ‘sharp’. f. The p orbital is a dumbbell shape. One major difference between s and p orbitals is that there are three of the p orbitals, all having identical energies. One p orbital lines up along the x axis, one along the y axis and one along the z axis. g. Each orbital has a maximum number of electrons it can hold: h. Electron configurations:i. This is an example of electron configurations. The number 1 in 1s indicates how far away from the nucleus the electron is. The letter s in 1s indicates the shape ofthe orbital. j. Example of sequencing: (same for all atoms) 1s, 2s, 2p, 3s, 3p etc…k. The difference between 1s and 2s is how far the electron is from the nucleus. 2s is a larger orbital than 1s and the electron in 2s will be in a more excited state. l. Electrons cannot be located between