Principles of Chemistry: Chapt 2. Atoms and Elements. Modern Atomic theory and the Laws that led to it. Law of Conservation of Mass

Transcription

1 Principles of Chemistry: A Molecular Approach, 1 st Ed. Nivaldo Tro Chapter 2 Atoms and Elements Modern Atomic theory and the Laws that led to it Atomic Theory- All matter is composed of atoms grew out of observation and laws Three important Laws that led to development and acceptance of Atomic Theory Law of conservation of mass Law of definite proportions Law of multiple proportions Law of Conservation of Mass In a chemical reaction, matter is neither created nor destroyed. (Antoine Lavoisier ) Total mass of the materials you have before the reaction must equal the total mass of the materials you have at the end. total mass of reactants = total mass of products 1

2 Reaction of Sodium with Chlorine to Make Sodium Chloride The mass of sodium and chlorine used is determined by the number of atoms that combine 7.7 g Na g Cl g NaCl Law of Definite Proportions All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements. Joseph Proust ( ) Show that two samples of carbon dioxide obey the Law of Definite Proportions Given: Sample 1: 25.6 g O and 9.60 g C Sample 2: 21.6 g O and 8.10 g C Find: Proportion O:C Conceptual g O1, g C 1 O:C in each sample Plan: g O2, g C 2 Relationships: All samples of a compound have the same proportion of elements by mass. Solution: Check: Since both samples have the same O:C ratio, the result is consistent with the Law of Definite Proportions. 2

3 Law of Definite Proportions Practice If a 10.0-g sample of calcite contains 4.0 g of calcium, how much calcite contains 0.24 g of calcium? Given:Sample 1: 4.0 g Ca and 10.0 g calcite Sample 2: 0.24 g Ca Find :mass calcite, g Law of Multiple Proportions When two elements A and B form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small whole numbers. (John Dalton ) For example oxides of carbon (CO and CO2) CO CO2 Ratio of these two masses itself is a small whole number Dalton s Atomic Theory 1) 2) 3) 4) Dalton explained these laws with his atomic theory which included the following concepts; Each element is composed of tiny, indestructible particles called atoms. All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements Atoms combine in simple, whole-number ratios to form compounds Atoms of one element cannot change into atoms of another element., In a chemical reaction atoms simply rearrange the way they are attached 3

4 The Discovery of electron Although Dalton postulated that atoms were indivisible, experiments at the beginning of this century showed that atom itself was composed of even smaller, more fundamental particles. J.J. Thompson In 1897 British physicist Cathode ray tube identified the electrons(subatomic particles) See Figure 2.3 calculated the ratio of the electron s mass, m e, to its electric charge, e (m/e) Cathode Ray Tube An evacuated glass tube containing metal electrodes when connected to a high-voltage power supply, produced a stream of particles called cathode rays. Properties of cathode rays Cathode rays- are emitted from negatively charged electrode, called cathode travel to positively charged electrode called anode Cathode rays are stream of negatively charged particles They traveled in straight lines were independent of the composition of the material from which they originated(cathode) 4

5 Some Notes on Charge opposite charges attract like charges repel To be neutral, something must have no charge or equal amounts of opposite charges. Thomson s Results The cathode rays are stream of negatively charged tiny particles because the beam always deflected toward the + plate The amount of deflection was related to two factors; the charge and mass of the particles Every material tested contained these same particles The charge:mass ratio of these particles was C/g 2000 times lighter than hydrogen the lightest atom known Millikan showed that the particle did have the same amount of charge as the hydrogen ion became known as electrons Millikan s Oil Drop Experiment In 1909, U.S. physicist, Robert Millikan had obtained the charge on the electron 5

6 Millikan s Oil Drop Experiment Oil droplets received charge from ionizing radiation Charged drops suspended in an electric field The mass and charge of each droplet were used to calculate the mass and charge of single electron Electrons Electrons are particles found in all atoms. Cathode rays are streams of electrons. The electron has a charge of C. The electron has a mass of g. Charge x mass /charge = mass C x g/-1.76 x 10 8 C = g. A New Theory of the Atom Thomson s new model of the atom replaced the first statement in Dalton s Atomic Theory. He proposed that atom actually has an inner structure atom is no longer indivisible Plum-Pudding Modle (J.J. Thompson) Nuclear Atom (Rutherford s Gold foil expt.) Thomson s Plum Pudding Atom The structure of the atom contains many negatively charged electrons held within a positively charged sphere (atom) There had to be a source of positive charge because the atom is neutral 6

7 Radioactivity Henri Becquerel and Marie Curie(1800), certain elements would constantly emit small, energetic particles and rays. These energetic particles could penetrate matter. three different kinds of emissions. alpha,, particles with a mass 4 H atom and + charge beta,, particles with a mass ~1/2000 H atom and charge gamma,, rays that are energy rays, not particles Structure of the atom In 1897, the British physicist J.J. Thompson showed that atoms were not indivisible particles. Thompson calculated the ratio of the electron s mass, m e, to its electric charge, e. In 1909, U.S. physicist, Robert Millikan had obtained the charge on the electron. These two discoveries combined provided us with the electron s mass of x kg, which is more than 2000 times smaller than the mass of the lightest atom (hydrogen). These experiments showed that the electron was indeed a subatomic particle Rutherford s Experiment Idea of the nuclear model of the atom based on experiments done in his laboratory by Hans Geiger and Ernest Morrison. (See Figure 2.5) Rutherford s famous gold Foil experiment majority of the atom s mass and all of its positive charge are contained in a small core called nucleus. Alpha particles directed at a thin gold film deflected in all directions including back at the source Only a large positive charge could cause the particles to bounce back 7

8 Rutherford s Nuclear Atom An atom is mostly empty space throughout which tiny negatively charged electrons are dispersed because almost all the particles went straight through Most of atoms s mass and all of its positive charge are contained in a small core called nucleus. because of the few particles that bounced back This dense particle was positively charged. There are as many negatively charged electrons outside the nucleus as there are positively charged particles(protons) within the the nucleus, so the atom is electrically neutral. A few of the particles do not go through. Plum Pudding Atom Nuclear Atom... If an atom was like a plum pudding, all the particles should go straight through. Almost all particles go straight through. Some particles(1 in ) go through, but are deflected due to +:+ repulsion from the nucleus. 8

9 Structure of the Atom Protons-has charge same as an electron but with opposite sign. Since protons and electrons have the same amount of charge, for the atom to be neutral there, must be equal numbers of protons and electrons. The nucleus of an atom is composed of two different kinds of particles, protons and neutrons. Subatomic particles Neutron ( Rutherford and Chadwick) have no charge a mass of 1 amu,slightly heavier than a proton Elements Atomic number (Z)-The number of protons in the nucleus of an atom Each element has a unique number of protons (Z) The elements are arranged on the Periodic Table in order of their atomic numbers. Each element has a unique name and symbol. symbol either one or two letters of the name in english, Greek or Latin one capital letter or one capital letter and one lowercase e.g. Hydrogen = H for sodium(natrium Latin)= Na Tin = Sn (stannum in Latin) 9

10 Atomic number Isotopes Isotopes- Atoms of the same element with same number of protons but different number of neutrons e.g for all carbon atoms Z= 6 (C-12 and C-13) The observed mass is a weighted average of the weights of all the naturally occurring atoms Natural abundance- The percentage of one isotope in an element The atomic mass of chlorine is amu Cl 17Cl 10

11 Isotopes All isotopes of an element are chemically identical. undergo the exact same chemical reactions have the same number of protons. have different masses due to different numbers of neutrons. Isotopes are identified by their mass numbers (A). Mass number A = protons (p) + neutrons (n) Isotopes A nuclide is an atom characterized by a definite atomic number and mass number. The shorthand notation for a sodium nuclide Atomic Number Z =Number of protons Mass Number A = Protons + Neutrons 23 sodium Mass number= neutron + proton Na Atomic number = # proton Neon Symbol Number of Protons Number of Neutrons A, Mass Number Percent Natural Abundance Ne-20 or 20 Ne % Ne-21 or Ne % Ne-22 or Ne % 11

12 Practice Complete the Table Protons Neutrons Electrons Atomic Number Mass Number Atomic Symbol Ions: losing and gaining electrons Ions- When atoms gain or lose electrons, they acquire a charge and are called ions Anions Negatively charged ions (atoms gain electrons) Cations - Positively charged ions (atoms lose electrons) Ions behave much differently than the neutral atom. Since materials like table salt are neutral, there must be equal amounts of charge from cations and anions in them. Atomic Structures of Ions Nonmetals form anions For each negative charge, the ion has 1 more electron than the neutral atom. F = 9 p + and 9 e -, F - = 9 p + and 10 e - Anions are named by changing the ending of the name to -ide fluorine F + 1e - F - fluoride ion oxygen O + 2e - O 2- oxide ion 12

13 IONS Metals form cations For each positive charge, the ion has 1 less electron than the neutral atom. Na atom = 11 p + and 11 e -, Na + ion = 11 p + and 10 e - Cations are named the same as the metal. sodium Na Na + + 1e - sodium ion The Periodic Law and Periodic Table Dimitri Mandeleev ordered elements by atomic mass Periodic Law When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically. put elements with similar properties in the same column used pattern to predict properties of undiscovered elements Where atomic mass order did not fit other properties, he reordered by other properties. Te & I Periodic Law 13

14 Periodic Table Periodic Table -Tabulates all known elements in order of increasing atomic number. A period consists of the elements in one horizontal row of the periodic table. A group consists of the elements in any one column of the periodic table. The groups are usually numbered. Main group The eight A groups are called main group (or representative) elements. (See Figure 2.11) The Periodic Table Transition elements The B groups are called transition elements. Inner transition elements. The two rows of elements at the bottom of the table Elements in any one group have similar properties. The elements in group IA, often known as the alkali metals, are soft metals that react easily with water. The group VIIA elements, known as the halogens, are also reactive elements Major Divisions of Periodic table Metals, nonmetals and metalloids A metal is a substance or mixture that has a characteristic luster and is generally a good conductor of heat and electricity. A nonmetal is an element that does not exhibit the characteristics of the metal. A metalloid, or semi-metal, is an element having both metallic and nonmetallic properties. 14

15 Patterns in Metallic Character = Metal = Metalloid = Nonmetal The Modern Periodic Table 15

16 Important Groups Hydrogen Nonmetal colorless, diatomic gas very low melting point and density reacts with nonmetals to form molecular compounds HCl is acidic gas H 2 O is a liquid reacts with metals to form hydrides metal hydrides react with water to form H 2 HX dissolves in water to form acids Alkali Metals Alkali metals: Group 1A metals Na, Li, K and Rb, soft, low melting points, low density, all are very reactive react violently with water to form basic (alkaline) solutions and H 2 (releases a lot of heat) 2 Na + 2 H 2 O 2 NaOH + H 2 Alkali Earth Metals: Group 2A, harder, higher melting, and denser than alkali metals less reactive than corresponding alkali metal form stable, insoluble oxides from which they are normally extracted alkaline earth oxides are basic reactivity with water to form H 2 Be = none; Mg = steam; Ca, Sr, Ba = cold water Halogens Halogens: Group 7A elements, nonmetals F 2 and Cl 2 gases; Br 2 liquid; I 2 solid all diatomic, very reactive Cl 2, Br 2 react slowly with water Br 2 + H 2 O HBr + HOBr react with metals to form ionic compounds HX all acids HF weak < HCl < HBr < HI 16

17 Noble Gases Noble Gases- Group 8A, all gases at room temperature very low melting and boiling points very unreactive, practically inert very hard to remove electron from or give an electron to Argon, a small component of the earth s atmosphere Ions and the Periodic Table The charge on an ion can often be determined from an element s position on the Periodic Table Metals always form positively charged cations For many main group metals, the charge = the group number. Nonmetals form negatively charged anions For main group nonmetals, the charge = the group number-8 Predictable Ion charges 17

18 What is the charge on each of the following ions? potassium cation sulfide anion calcium cation bromide anion aluminum cation K + S 2- Ca 2+ Br - Al 3+ Atomic Mass We generally use the average mass of all an element s atoms found in a sample in calculations However, the average must take into account the abundance of each isotope in the sample Atomic mass- The average mass of an elements atoms At. Wt. = m A x(its fractional abundance)+ m B x (its fractional abundance) Example 2.5 If copper is 69.17% Cu 63 with a mass of amu and the rest Cu 65 with a mass of amu, find copper s atomic mass Given: Cu 63 = 69.17%, amu Cu 65 = %, amu Find: atomic mass, amu Conceptual Plan: Relationships: isotope masses, isotope fractions avg. atomic mass Solution: Check: The average is between the two masses, closer to the major isotope. 18

19 Counting Atoms by Moles: Mole (A chemist s dozen): One mole of an element is the amount of that element that contains Avogadro s number(6.022x ) of atoms. 1 mole = units like 1 dozen = 12 things One mole is equal to the number of atoms is equal to the number of atoms in 12 g of pure C-12 1 atom of C-12 weighs exactly 12 amu. 1 mole of C-12 weighs exactly 12 g. 1 mole of C atoms =12.01 g = C atoms Example 2.6 Calculate the number of atoms in 2.45 mol of copper Given: Find: Conceptual Plan: Relationships: Solution: mol Cu 2.45 mol Cu atoms Cu atoms Cu 1 mol = atoms Check: Since atoms are small, the large number of atoms makes sense. Converting between Mass and amount (Moles) Molar mass-the mass of one mole of atoms of an element The molar mass of an element, in grams, is numerically equal to the element s atomic mass, in amu. 1 mole O = g O = amu = 6.022x mol C = g C = 6.022x O atoms Conversion factor 1 mol O/16.0 g O or 16.0 g O/ 1 mol O O atoms 19

20 Sample Problem- Mass to moles Calculate the moles of carbon in g of pencil lead. Learning objectives To know Modern Atomic theory Using Law of definite proportions and multiple proportions Working with atomic numbers, mass numbers and isotope symbols Determining atomic weight from isotopic masses and fractional abundances Predicting the charge of the ions Calculating atomic mass, converting between mole and number of atoms Using mole concept and converting between mass and moles and us 20

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