Chem Chapter 3 PDF

Chapter 3: Atoms – The Building Blocks of Matter Section 1: The Atom – From Philosophical Idea to Scientific Theory Early philosophers pondered the fundamental nature of matter. Is it continuous and infinitely divisible, or is it divisible only until a basic, invisible particle that c...

Chapter 3: Atoms – The Building Blocks of Matter Section 1: The Atom – From Philosophical Idea to Scientific Theory Early philosophers pondered the fundamental nature of matter. Is it continuous and infinitely divisible, or is it divisible only until a basic, invisible particle that cannot be divided further is reached? The particle theory of matter was supported as early as 400 BCE by the Greek philosopher Democritus. o He called nature’s basic particle an atom, based on the Greek word atomos, meaning indivisible. Aristotle was part of the generation that succeeded Democritus. He did not believe in atoms; he thought matter was continuous. His opinion was accepted for nearly 2000 years. Neither Democritus nor Aristotle had experimental evidence to support their claims, so each remained under speculation until the eighteenth century. Three basic laws describe how matter behaves in chemical reactions. Virtually all chemists in the late 1700s accepted the modern definition of an element as a substance that cannot be broken down further by ordinary chemical means. They also assumed that elements combined to form compounds that have different physical and chemical properties. What they did not understand was exactly how the different substances could combine with one another to form new substances, what we know as chemical reactions. The late1790s 1970s saw the improvement of balances, which allowed scientists to accurately measure the masses of elements and compounds they were studying. This led to the discovery of several basic laws. The Law of Conservation of Mass: matter is neither created nor destroyed during chemical reactions or s physical changes. The Law of Definite Proportions: the fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the sample size or source of the compound. o Sodium chloride (table salt) always consists of 39.34% by mass of sodium and 60.66% by mass of chloride, no matter how much sample is measured. The Law of Multiple Proportions: If two or more different compounds are composed of the same two elements , then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers. o The elements carbon and oxygen form the compounds carbon dioxide and carbon monoxide. Consider an example where in both compounds, there is 1.00 gram of carbon. In carbon dioxide, Always 2.66 grams of oxygen combine with 1.00 gram of carbon. In carbon monoxide, 1.33 grams of oxygen combines with 1.00 gram of carbon. The ratio of the masses of oxygen is 2.66 to 1.33, followthis or 2 to 1. rule know definition as 1.33325W  Compounds contain atoms in whole-number ratios. In 1808, an English school teacher named John Dalton proposed an explanation that encompassed all of these laws. His theory can be summed up by the following statements. Dalton’s atomic theory: 1. All matter is composed of extremely small particles called atoms. 2. Atoms of an element are identical in size, mass, and other properties; atoms of different elements not do not have identical traits. identical in 3. Atoms cannot be subdivided, created, or destroyed. Msn be subdived 4. Atoms of different elements combine in simple whole number ratios to form chemical compounds 5. In chemical reactions, atoms are combined, separated, or rearranged. Dalton’s atomic theory explains the law of conservation of mass through the concept that chemical reactions involve merely the combination, separation, or rearrangement of atoms and that during reactions atoms are not subdivided, created, or destroyed. Atoms can be subdivided into smaller particles. Not all aspects of Dalton’s atomic theory have proven to be correct. Today we know that atoms are divisible into even smaller particles (although the law of conservation of mass still holds true for chemical reactions). We also know that a given element can have atoms with different masses. The important concepts that (1) all matter is composed of atoms and (2) atoms of any one element differ in properties from atoms of another element remain unchanged. timeline Democritus B C 1800s Dalton 1700s early Thomson Late 1800s cathoderarefactions Milikan hmassive hothead Rubato Nucleus ilExp  electron cloud Chapter 3: Atoms – The Building Blocks of Matter Section 2: The Structure of the Atom Although John Dalton thought atoms were indivisible, investigators in the late 1800s proved otherwise. It became clear that atoms are actually composed of smaller particles and that the number and arrangement of these particles within an atom determine the atom’s chemical properties. Therefore, today we define an atom as the smallest particle of an element that retains the chemical properties of that element. Atoms consist of two regions: 1. The nucleus: a very small region located at the center of an atom. § The nucleus of any atom contains at least one positively charged particle called a proton, and usually one or more neutral particles called neutrons. 2. Surrounding the nucleus is a region occupied by negatively charged particles called electrons. Protons, neutrons, and electrons are referred to as subatomic particles. Atoms contain positive and negative particles. In the late 1800s, many experiments were performed in which electric current was passed through various gases at low pressures. (Gases at normal atmospheric pressure do not conduct electricity well.) These experiments were carried out in glass tubes, called cathode-ray tubes, that had been hooked up to a vacuum pump. Cathode Rays and Electrons Investigators noticed that when current was passed through the tube, the surface of the tube directly opposite the cathode glowed. They hypothesized that the glow was caused by a stream of particles, which they called the cathode ray. The ray traveled from the cathode to the anode when current was passed through the tube. The following observations were made: 1. Cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have a negative charge. 2. The rays were deflected away from a negatively charged object. These observations led to the hypothesis that the particles that compose cathode rays are negatively charged. In 1897, English physicist J.J. Thomson was able to measure the ratio of the charge of cathode ray particles to their mass. He found that this ratio was always the same, regardless of the metal that was used to make the cathode or the nature of the gas inside the cathode ray tube. Thomson concluded that all cathode rays must be composed of negatively charged particles, which were named electrons. Koformat  Charge and Mass of the Electron Cathode rays have identical properties, regardless of the element used to produce them. Therefore, it was concluded that electrons are present in atoms of all elements. This provided evidence that atoms are divisible and that one of the atoms basic constituents is the negatively charged electron. Didmathfor JJ In 1909, physicist Robert Millikan measured the charge of Thompson's expirement the electron, and scientists were able to use this information to determine that the mass of an electron is about 2000X smaller than a hydrogen atom (the smallest of all atoms) The mass of an electron is 9.109 x 10-31 kg Based on what was learned about electrons, scientists inferred two other facts about the atom: 1. Because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons. 2. Because electrons have such a small mass, atoms must contain other particles that account for most of their mass. Thomson proposed a model for the atom called the plum pudding model (after the English dessert). a o He believed that the negative electrons were spread evenly throughout the positive charge of the rest of the atom. o Although later experiments disproved it, Thomson’s plum pudding model represents the first time scientists tried to incorporate the idea that atoms were not indivisible. Atoms have small, dense, positively charged nuclei. In 1911, Ernest Rutherford performed an experiment where a thin piece of gold foil was bombarded with fast moving alpha particles, which are positively charged particles about 4X bigger than a hydrogen atom. With the plum pudding atomic model in mind, the expectation was for the alpha particles to pass through the foil with very little deflection. For the vast majority of particles, this was the case. But it was found that a small percentage of the alpha particles were deflected back in the direction of the source. o “It was as if you fired a 15 inch bullet at a piece of tissue paper and it came back and hit you” - Rutherford Rutherford concluded that atoms must contain a very small, densely packed bundle of matter with a positive electric charge. He named this positive bundle the nucleus. positive charge is than the more negative charge in atoms A nucleus contains protons and neutrons. A proton has a positive charge that is equal in magnitude to the negative charge of an electron. Atoms are electrically neutral because they contain equal numbers of protons and electrons. Neutrons are electrically neutral (no charge). The mass of a proton is 1.673 x 10-27 kg. The mass of a neutron is 1.675 x 10-27 kg. Atoms of different elements differ in their number of protons and, therefore, the amount of positive charge they possess. o The number of protons determines an atoms identity (what element it is). Forces in the Nucleus Generally, particles that have the same electric charge repel one another. So, you’d expect a nucleus with more than one proton to be unstable. However, when two protons are extremely close to each other, there is a strong attraction between them. This also occurs when neutrons and protons are closely together. Nuclear forces are short range proton-neutron, proton-proton, and neutron-neutron forces that hold the nucleus together. The radii of atoms are expressed in picometers. The region occupied by electrons surrounding the nucleus is referred to as the electron cloud. The atomic radius is the distance from the center of the nucleus to the outer portion of the electron cloud. o Think of atomic radius as the size of the atom. Because atomic radii are so small, they are expressed in picometers. (1pm = 10-12m)  Chapter 3: Atoms – The Building Blocks of Matter Section 3: Counting Atoms All atoms of an element must have the same number of protons but not neutrons. The atomic number of an element is the number of protons in each atom of that element. The atomic number identifies an element. If the number of protons in the nucleus of an atom were to change, that atom would become a different element. Isotopes Just because all hydrogen atoms, for example, have 1 proton, it doesn’t mean they all have the same number of neutrons. In fact, three types of hydrogen atoms are known. hydrogen 1. Protium: the most common type of hydrogen (99.99% abundance); it contains 1p+ , 0 n0 2. Deuterium: accounts for 0.01% of hydrogen; contains 1p+, 1n0 + 0 sotopes 3. Tritium: manmade and radioactive. Not very common. Contains 1p , 2n Protium, deuterium, and tritium are isotopes of hydrogen. Isotopes are atoms of the same element that have different masses. o Isotopes have the same number of protons and electrons, but different numbers of neutrons. Although isotopes have different masses, they do not differ significantly in their chemical behavior. Mass Number Identifying an isotope requires knowing both the name or atomic number of the element and the mass of the isotope. The mass number is the total number of protons and neutrons that make up the nucleus of an isotope. Identifying Isotopes Generalrelativesense of its weight There are two methods for specifying isotopes: 1. Hyphen notation: write the name of the element, hyphen, then mass number. For example, Tritium is written Hydrogen – 3 2. Nuclear symbol: Superscript the mass number, subscript the atomic number, then element symbol. For example, Tritium is written 3H 1 The number of neutrons is found by subtracting the atomic number from the mass number. o Number of neutrons in Tritium: Mass number (3) – atomic number (1) = 2 neutrons Nuclide is a general term for a specific isotope of an element. Calculating the Actual Mass of an Isotope To calculate the mass of a specific isotope: (number of p+ X mass of proton) + (number of n0 X mass of neutron) + (number of e- X mass of electron) Atomic mass is a relative measure. Masses of atoms expressed in grams are very small. For most chemical calculations, it is more convenient to use relative atomic masses. The standard used by scientists to compare units of atomic mass is the carbon-12 atom, which has been arbitrarily assigned a mass of exactly 12 atomic mass units, or 12 amu. One atomic mass unit (amu) is exactly 1/12 the mass of a carbon-12 atom. The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom. o Example: The hydrogen-1 atom has a mass of 1.01 amu because it is about 1/12 the mass of a carbon-12 atom. The oxygen-16 atom has a mass of about 16/12 the mass of a carbon-12 atom. Average atomic mass is a weighted value. Most elements occur naturally as mixtures of isotopes. Scientists determine the average mass of a sample of an element’s isotopes by determining the percentages of each of the isotopes then giving the proper weight to each value. Average Atomic Mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. Calculating average atomic mass: multiply the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results. Round atomic mass to 2 decimal places before calculating. Example: Isotope Atomic Mass (amu) Relative abundance (%) Copper - 63 62.929601 69.15 Copper - 65 64.927794 30.85 (62.929601 amu X 0.6915) + (64.927794 X 0.3085) = 63.55 amu A relative mass scale makes counting atoms possible. 2ⁿᵈ decimalplace The relative atomic mass scale makes it possible to know how many atoms of an element are present in a sample. Three very important concepts: the mole, Avogadro’s number, and molar mass, provide the basis for relating masses in grams to number of atoms. The Mole A mole (mol) is the amount of substance that contains as many particles as there are atoms in exactly 12 grams of carbon-12. The mole is a counting unit, just like a dozen is. We don’t usually order 12 or 24 donuts; we order 1 or 2 dozen. Similarly, a chemist may want 1 mol of carbon, or 2 mol of iron, or 2.567 mol of calcium. Avogadro’s Number Avogadro’s number (6.022 x 1023) is the number of particles (atoms or molecules) in exactly one mole of a pure substance. Molar Mass Molar mass is the mass of one mole of a pure substance. Molar mass is usually written in units of g/mol. The molar mass of an element is equal to the atomic mass of an element found on the periodic table. 1 mole of any element/compound = 6.022 x 1023 atoms/molecules = the Molar Mass of element/compound 7.5g Nad 1 mol 0.13 58.43g 1m01 6.022 107 8 9 organic chemistry tutor youtube mass  Patin's mols Test Format Multiple choice 9 Fill in the blank 10 Discussion 2 - 15 points Isotope chart 3, ll in the blank - 6 pts Calculations 4, 23 points