Presentation on theme: "Atoms and Elements Chapter 4 Tro, 2 nd ed.. All known substances on Earth and probably the universe are formed by combinations of more than 100 elements."— Presentation transcript:
All known substances on Earth and probably the universe are formed by combinations of more than 100 elements. An element is a fundamental or elementary substance that cannot be broken down into simpler substances by chemical means. Each element has a number. Beginning with hydrogen as 1, the elements are numbered in order of increasing complexity. ELEMENTS AND ATOMS
Most substances can be decomposed into two or more simpler substances. Water can be decomposed into hydrogen and oxygen. Table salt can be decomposed into sodium and chlorine. An element cannot be decomposed into a simpler substance. An atom is the smallest particle of an element that can exist. An atom is the smallest unit of an element that can enter into a chemical reaction. ELEMENTS AND ATOMS
The diameter of an atom is 0.1 to 0.5 nm. This is 1 to 5 ten billionths of a meter. If the diameter of this dot is 1 mm, then 10 million hydrogen atoms would form a line across the dot. Even smaller particles than atoms exist. These are called subatomic particles.
2.Atoms of the same element are alike in mass and size. Atoms of different elements have different masses and sizes. (Is this still correct?) 3.Chemical compounds are formed by the union of two or atoms of different elements in whole number ratios. Dalton’s Atomic Theory Modern research has demonstrated that atoms are composed of subatomic particles. Atoms under special circumstances can be decomposed. 1. Elements are composed of minute indivisible particles called atoms.
In 1897 Sir Joseph Thompson demonstrated that cathode rays: - travel in straight lines - are negative in charge - are deflected by electric and magnetic fields - produce sharp shadows - are capable of moving a small paddle wheel WHAT’S IN AN ATOM? DISCOVERING SUBATOMIC PARTICLES!
This was the discovery of the fundamental unit of charge – the electron. Subatomic Particles: Electron
Eugen Goldstein, a German physicist, first observed protons in 1886: Thompson determined the proton’s characteristics. Thompson showed that atoms contained both positive and negative charges. This disproved the Dalton model of the atom which held that atoms were indivisible. But the mass of the atom could not be accounted for by the mass of protons inside it. There had to be something else. Subatomic Particles: Protons
James Chadwick discovered the neutron in 1932. Its actual mass is slightly greater than the mass of a proton. Subatomic Particles: Neutron
Memorize name, symbol, relative charge, relative mass, and LOCATION (in nucleus or in orbitals outside of nucleus) of these three subatomic particles. See table 4.1 on page 90/96 in your book. 4.1 0.00055 1.0073 1.0087
Radioactivity was discovered by Becquerel in 1896. Radioactive elements spontaneously emit alpha ( ) particles, beta ( ) particles and gamma ( ) rays from their nuclei. By 1907, Rutherford found that particles emitted by certain radioactive elements were helium nuclei, consisting of 2 protons and 2 neutrons. The Nuclear Atom
In 1911, Rutherford performed experiments that shot a stream of particles at a thin piece of gold foil. Most of the particles passed through the foil with little or no deflection. He found that a few were deflected at large angles and some particles even bounced back. The Rutherford Experiment
An electron with a mass of 0.00055 amu could not have deflected an particle with a mass of ~4 amu. Rutherford knew that like charges repel. Rutherford concluded that each gold atom contained a positively charged mass that occupied a tiny volume. He called this mass the nucleus. The Rutherford Experiment
If a positive particle approached close enough to the positive mass it was deflected. Most of the particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space. Because particles have relatively high masses, the extent of the deflections led Rutherford to conclude that the nucleus was very heavy and dense. The Rutherford Experiment
5.5 Deflection and scattering of particles by positive gold nuclei. Deflection Scattering
Rutherford’s experiment showed that an atom had a dense, positively charged nucleus. Chadwick’s work in 1932 demonstrated the atom contains neutrons. Rutherford also noted that light, negatively charged electrons were present in an atom and offset the positive nuclear charge. General Arrangement of Subatomic Particles
Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center. The negative electrons surround the nucleus. The nucleus contains protons and neutrons General Arrangement of Subatomic Particles
The atomic number of an element is equal to the number of protons in the nucleus of that element. The atomic number of an atom determines which element the atom is. Every atom with an atomic number of 1 is a hydrogen atom. Every carbon atom contains 6 protons in its nucleus. ATOMS AND THEIR SUBATOMIC PARTICLES
1 proton in the nucleus atomic number Every atom with an atomic number of 1 is a hydrogen atom. 1H1H
6 protons in the nucleus 6C6C atomic number Every atom with an atomic number of 6 is a carbon atom.
Positive ions were explained by assuming that a neutral atom loses electrons. Negative ions were explained by assuming that extra electrons can be added to atoms. Ions: atoms that have lost or gained electrons
5.4 When one or more electrons are lost from an atom, a cation is formed.
5.4 When one or more electrons are added to a neutral atom, an anion is formed.
Atoms of the same element can have different masses. They always have the same number of protons, but they can have different numbers of neutrons in their nuclei. The difference in the number of neutrons accounts for the difference in mass. These are isotopes of the same element. Isotopes of the Elements
The mass of a single atom is too small to measure on a balance. Using a mass spectrometer, the mass of the hydrogen atom was determined. The mass of one hydrogen atom was determined to be 1.673 x 10 -24 g. Atomic Mass
A Modern Mass Spectrometer A mass spectrogram is recorded. From the intensity and positions of the lines on the mass spectrogram, the different isotopes and their relative amounts can be determined. Positive ions formed from sample. Electrical field at slits accelerates positive ions. Deflection of positive ions occurs at magnetic field. 5.8
5.9 A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given.
Single atomic masses are too small to weigh on a balance. To overcome this problem a system of relative atomic masses using “atomic mass units” was devised to express the masses of elements using simple numbers. The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon, carbon-12. A mass of exactly 12 atomic mass units (amu) was assigned to the carbon- 12 atom. An amu is defined as exactly equal to 1/12 th mass of a carbon-12 atom. 1 amu = 1.6606 x 10 -24 g Isotopes of the same element have different masses. The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12. ATOMIC MASS UNITS AND RELATIVE ATOMIC MASS
Calculating relative atomic mass: Convert percent abundance to fraction by dividing by 100 Multiply fraction abundance times each isotopes mass Sum up the results Rel Atomic Mass = f 1 *m 1 + f 2 *m 2 + … See example on next slide
To calculate the atomic mass multiply the atomic mass of each isotope by its fractional abundance and add the results. (62.9998 amu) 0.6909 =43.48 amu (64.9278 amu)0.3091 =20.07 amu 63.55 amu Isotope Isotopic mass (amu) Abundance (%) Average atomic mass (amu) 62.999869.09 64.927830.91 63.55
Relationship between Mass Number and Atomic Mass Atomic mass found on the Periodic Table is RELATIVE atomic mass, the weighted average of all the isotopes. Each isotope of an element has a unique WHOLE number called the mass number, the sum of protons and neutrons. Mass number is NOT found on most Periodic Tables. ****************************************************************** ATOMIC MASS IS NOT THE SAME AS MASS NUMBER ****************************************************************** REMEMBER: ATOMIC MASS IS AN AVERAGE OF THE DIFFERENT ISOTOPES’ MASSES. It will have decimal points!
The mass number minus the atomic number equals the number of neutrons in the nucleus. mass number atomic number number of neutrons - = 62109-47= mass number atomic number Relationship Between Mass Number and Atomic Number
Atomic Mass, Mass Number & Abundance of Isotopes EXAMPLE: Chlorine has two isotopes, chlorine-35 and chlorine-37. Write down the atomic number, mass number and number of neutrons for each. Then look up the atomic mass. Which isotope is in greater abundance? Cl-35: Z = 17, A = 35, 35 - 17 = 18 neutrons Cl-37: Z = 17, A = 37, 37 - 17 = 20 neutrons Atomic mass is 35.453 amu. Closer to 35, therefore chlorine-35 is more abundant.
Elements are not distributed equally by nature. Oxygen is the most abundant element in the human body (65%). Oxygen is the most abundant element in the crust of the earth (49.2%). In the universe, the most abundant element is hydrogen (91%) and the second most abundant element is helium (8.75%).
3.2 Distribution of the common elements in nature.
Sources of Element Names Famous- Scientists Einsteinium: named for Albert Einstein. LocationGermanium: discovered in 1866 by a German chemist. German- Color Bismuth: from the German weisse mass which means white mass. Greek- Color Iodine: from the Greek iodes meaning violet. Latin- Property Fluorine: from the Latin fluere meaning to flow. The fluorine containing ore fluorospar is low melting.
Most symbols start with the same letter as the element.A number of symbols appear to have no connection with the element. These symbols have carried over from the earlier names of the elements (usually Latin).
Element Names and Symbols Element names and symbols are in Periodic Table inside front cover. Symbols: Learn first 36 on Periodic Table, plus Ag, Au, Pt, Hg, Sn, Pb, I, U, Ba, and Rn. Also memorize that some elements exist as diatomic molecules : H 2, O 2, N 2, F 2, Cl 2, Br 2, and I 2
The periodic table was designed by Dimitri Mendelev in 1869. In the table each element’s symbol is placed inside of a box. Above the symbol of the element is its atomic number. N 7 THE PERIODIC TABLE OF THE ELEMENTS
The elements are arranged in order of increasing atomic number. Elements with similar chemical properties are organized in columns called families or groups. HeKrNeArXeRn These elements are known as the noble gases. They are nonreactive.
THE PERIODIC TABLE OF THE ELEMENTS 18 columns - called groups 7 rows - called periods Representative elements are in groups 1,2, 3A-8A (13-18) Transition metals are in groups 3-12 Groups of elements have similar chemical properties
THE PERIODIC TABLE OF THE ELEMENTS SPECIAL GROUP NAMES: 1A (1) = alkali metals 2A (2) = alkaline earth metals 8A (18) = noble gases (six) 7A (17) = halogens (four) NOTE:H really belongs to its own group, which is why it’s shown by itself in my periodic table! It has 1 electron in outer shell, like the 1A elements, but it’s not a metal, and reacts more like group 7A.
The Periodic Table of Elements Nonmetals: upper right Metals: lower left Metalloids touch stairs, except for Al.
Most elements are metals Metals are solid at room temperature. –Mercury is an exception. At room temperature it is a liquid. Metals have high luster (they are shiny). Metals are good conductors of heat and electricity. Metals are malleable (they can be rolled or hammered into sheets). physical properties of metals
Most elements are metals Metals are ductile (they can be drawn into wires). Most metals have a high melting point. Metals have high densities
Many metals readily combine with nonmetals to form ionic compounds. They can combine with sulfur. Metals have little tendency to combine with each other to form compounds. chlorine. In nature, minerals are formed by combinations of the more reactive metals with other elements. Chemical Properties of Metals oxygen.
A few of the less reactive metals such as copper, silver and gold are found in the free state. Metals can mix with each other to form alloys. Brass is a mixture of copper and zinc. Bronze is a mixture of copper and tin. Steel is a mixture of carbon and iron. Chemical Properties of Metals
Lack luster (they are dull) Have relatively low melting points Have low densities. Poor conductors of heat and electricity At room temperature, carbon, phosphorous, sulfur, selenium, and iodine are solids. Physical Properties of Nonmetals Everything metals are, nonmetals are NOT!
Solid sulfur selenium Physical State at Room Temperature phosphorouscarbon iodine NONMETALS
liquid Physical State at Room Temperature bromine NONMETALS
gas Physical State at Room Temperature helium, neon, argon, krypton, xenon, radon nitrogen, oxygen fluorine, chlorine NONMETALS
The Metalloids 1. boron 2. silicon 3. germanium 4. arsenic 5. antimony 6. tellurium 7. polonium Metalloids have properties that are intermediate between metals and nonmetals
Important Groups - Hydrogen Nonmetallic Colorless, diatomic gas – exists as H 2 Very low melting point & 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
65 Important Groups – IA, Alkali Metals Hydrogen usually placed here, though it doesn’t belong Soft, low melting points,low density Flame tests Li = red, Na = yellow, K = violet Very reactive, never find uncombined in nature Tend to form water soluble compounds with colorless solutions React with water to form basic (alkaline) solutions and H 2 2 Na + 2 H 2 O 2 NaOH + H 2 releases a lot of heat lithium sodium potassium rubidium cesium
Important Groups – IIA, Alkali Earth Metals Harder, higher melting, and denser than alkali metals Flame tests Ca = red, Sr = red, Ba = yellow-green Reactive, but less than corresponding alkali metal Form stable, insoluble oxides from which they are normally extracted Oxides are basic = alkaline earth Reactivity with water to form H 2, Be = none; Mg = steam; Ca, Sr, Ba = cold water magnesium calcium beryllium strontium barium
Important Groups – VIIA, Halogens Nonmetals F 2 & 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 bromine iodine chlorine fluorine
Important Groups – VIIIA, Noble Gases 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
Review What is the atomic number of boron, B? What is the atomic mass of silicon, Si? How many protons does a chlorine atom have? How many electrons does a neutral neon atom have? Will an atom with 6 protons, 6 neutrons and 6 electrons be electrically neutral? Will an atom with 27 protons, 32 neutrons and 27 electrons be electrically neutral? Will a Na atom with 10 electrons be electrically neutral? What is the mass number of fluorine?