{"id":2032,"date":"2023-10-02T08:36:45","date_gmt":"2023-10-02T08:36:45","guid":{"rendered":"http:\/\/localhost\/ecole9ja\/?p=2032"},"modified":"2023-10-02T08:43:01","modified_gmt":"2023-10-02T08:43:01","slug":"week-5-ss1-first-term-chemistry-notes-2","status":"publish","type":"post","link":"https:\/\/ecolebooks.com\/nigeria\/posts\/week-5-ss1-first-term-chemistry-notes-2\/","title":{"rendered":"Week 5 – SS1 First Term Chemistry Notes"},"content":{"rendered":"

WEEK 5
\n<\/strong>PARTICULATE NATURE OF MATTER<\/p>\n

\u00a0\"\"\/\"\"\/
\n\t\t\"\"\/
\n\t\tAtoms are the basic units of matter and the defining structure of elements. Atoms are made up of three particles: protons, neutrons and electrons.
\nProtons and neutrons are heavier than electrons and reside in the center of the atom, which is called the nucleus. Electrons are extremely lightweight and exist in a cloud orbiting the nucleus. The electron cloud has a radius 10,000 times greater than the nucleus.
\nProtons and neutrons have approximately the same mass. However, one proton weighs more than 1,800 electrons. Atoms always have an equal number of protons and electrons, and the number of protons and neutrons is usually the same as well. Adding a proton to an atom makes a new element, while adding a neutron makes an isotope, or heavier version, of that atom. <\/p>\n

Nucleus
\n<\/h2>\n

The nucleus was discovered in 1911, but its parts were not identified until 1932. Virtually all the mass of the atom resides in the nucleus. The nucleus is held together by the “strong force,” one of the four basic forces in nature. This force between the protons and neutrons overcomes the repulsive electrical force that would, according to the rules of electricity, push the protons apart otherwise.<\/p>\n

Protons
\n<\/h2>\n

Protons are positively charged particles found within atomic nuclei. They were discovered by Ernest Rutherford in experiments conducted between 1911 and 1919.
\nThe number of protons in an atom defines what element it is. For example, carbon atoms have six protons, hydrogen atoms have one and oxygen atoms have eight. The number of protons in an atom is referred to as the atomic number of that element. The number of protons in an atom also determines the chemical behavior of the element. The Periodic Table of the Elements arranges elements in order of increasing atomic number.
\nProtons are made of other particles called quarks. There are three quarks in each proton \u2014 two “up” quarks and one “down” quark \u2014 and they are held together by other particles called gluons.<\/p>\n

Electrons
\n<\/h2>\n

Electrons have a negative charge and are electrically attracted to the positively charged protons. Electrons surround the atomic nucleus in pathways called orbitals.\u00a0The inner orbitals surrounding the atom are spherical but the outer orbitals are much more complicated.
\nAn atom’s electron configuration is the orbital description of the locations of the electrons in an unexcited atom. Using the electron configuration and principles of physics, chemists can predict an atom’s properties, such as stability, boiling point and conductivity.
\nTypically, only the outermost electron shells matter in chemistry. The inner electron shell notation is often truncated by replacing the long-hand orbital description with the symbol for a noble gas in brackets. This method of notation vastly simplifies the description for large molecules.
\nFor example, the electron configuration for beryllium (Be) is 1s2<\/sup>2s2<\/sup>, but it’s is written [He]2s2<\/sup>. [He] is equivalent to all the electron orbitals in a helium atom. The Letters, s, p, d, and f designate the shape of the orbitals and the superscript gives the number of electrons in that orbital.<\/p>\n

Neutrons
\n<\/h2>\n

Neutrons are uncharged particles found within atomic nuclei. A neutron’s mass is slightly larger than that of a proton. Like protons, neutrons are also made of quarks \u2014 one “up” quark and two “down” quarks. Neutrons were discovered by James Chadwick in 1932.<\/p>\n

\u00a0
\n\u00a0RULES FOR FILLING ORBITALS<\/strong>
\n\t\tRule 1 <\/strong>– Lowest energy orbitals fill first. Thus, the filling pattern is 1s<\/strong>, 2s<\/strong>, 2p<\/strong>, 3s<\/strong>, 3p<\/strong>, 4s<\/strong>, 3d<\/strong>, etc. Since the orbitals within a subshell are degenerate (of equal energy), the entire subshell of a particular orbital type is filled before moving to the next subshell of higher energy.
\nRule 2<\/strong> – Pauli Exclusion Principle – Only two electrons are permitted per orbital and they must be of opposite spin. If one electron within an orbital possesses a clockwise spin, then the second electron within that orbital will possess a counterclockwise spin. Two electrons with opposite spins found in the same orbital are referred to as being paired.
\nRule 3<\/strong>– Hund’s Rule – The most stable arrangement of electrons in a subshell occurs when the maximum number of unpaired electrons exist, all possessing the same spin direction. This occurs due to the degeneracy of the orbitals, all orbitals within a subshell are of equal energy. Electrons are repulsive to one another and only pair after all of the orbitals have been singly filled. <\/p>\n

Rules for Assigning Electron Orbitals
\n<\/h3>\n

Occupation of Orbitals
\n\t\t\t\t<\/em><\/h4>\n

Electrons fill orbitals in a way to minimize the energy of the atom. Therefore, the electrons in an atom\u00a0fill the principal energy levels in order of increasing energy (the electrons are getting farther from the nucleus). The order of levels filled looks like this:\u00a0
\n1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p<\/em>
\n\t\tOne way to remember this pattern, probably the easiest, is to refer to the periodic table and remember where each orbital block falls to logically deduce this pattern. Another way is to make a table like the one below and use vertical lines to determine which subshells correspond with each other.\u00a0
\n\"\"\/<\/p>\n

Pauli Exclusion Principle
\n\t\t\t\t<\/em><\/h4>\n

The\u00a0Pauli\u00a0exclusion principle\u00a0states that no two electrons can have the same four quantum numbers. The first three (n,\u00a0l, and ml<\/sub>) may be the same, but the fourth quantum number must be different. A single orbital can hold a maximum of two electrons, which must <\/strong>have opposing spins; otherwise they would\u00a0have the same four quantum numbers, which is forbidden. One electron is spin up (ms<\/sub> = +1\/2) and the other would spin down (ms<\/sub> = -1\/2). This tells us that each subshell has double the electrons per orbital. The s subshell has 1 orbital that can hold up\u00a0to 2 electrons, the p subshell has 3 orbitals that can hold up to 6 electrons, the d subshell has 5 orbitals that hold up to 10 electrons, and the f subshell has 7 orbitals with 14 electrons.
\nExample 1: Hydrogen and Helium
\nThe first three quantum numbers of an electron are n=1, l=0, ml<\/sub>=0. Only two electrons can correspond to these, which would be either ms<\/sub> = -1\/2\u00a0or ms<\/sub> = +1\/2. As we already know from our studies of quantum numbers and electron orbitals, we can conclude that these four quantum numbers refer to the 1s subshell. If only one of the ms<\/sub>\u00a0values are given then we would have 1s1<\/sup>\u00a0(denoting hydrogen) if both are given we would have 1s2\u00a0<\/sup>(denoting helium). Visually, this is be represented as:<\/p>\n

Hund’s Rule
\n\t\t\t\t<\/em><\/h4>\n

When assigning electrons in orbitals, each electron will first fill all the orbitals with similar energy (also referred to as degenerate) before pairing with another electron in a half-filled orbital. Atoms at ground states tend to have as many unpaired electrons as possible. When visualizing this processes, think about how electrons are exhibiting the same behavior as the same poles on a magnet would if they came into contact; as the negatively charged electrons fill orbitals they first try to get as far as possible from each other before having to pair up.
\nExample 2: Oxygen and Nitrogen
\nIf we look at the correct electron configuration of the Nitrogen (Z = 7) atom, a very important element in the biology of plants: 1s2<\/sup> 2s2<\/sup> 2p3<\/sup>
\n\t\t\"\"\/
\n\t\tWe can clearly see that p orbitals are half-filled as there are three electrons and three p orbitals. This is because Hund’s Rule states that the three electrons in the 2p subshell will fill all the empty orbitals first before filling orbitals with electrons in them. If we look at the element after Nitrogen in the same period, Oxygen (Z = 8) its electron configuration is: 1s2<\/sup> 2s2<\/sup> 2p4<\/sup> (for an atom).
\n\"\"\/
\n\t\tOxygen has one more electron than Nitrogen and as the orbitals are all half filled the electron must pair up.
\n\u00a0
\n\u00a0<\/p>\n

The Aufbau Process
\n\t\t\t\t<\/em><\/h4>\n

Aufbau comes from the German word “aufbauen” meaning “to build.”\u00a0When writing electron configurations,\u00a0orbitals are built up from atom to atom. When writing\u00a0the electron configuration for an atom,\u00a0orbitals are filled in order of increasing atomic number.\u00a0However, there are some exceptions to this rule.
\nExample 3: 3rd<\/sup> row elements
\nFollowing the pattern across a period from B (Z=5) to Ne (Z=10), the number of electrons increases and the subshells are filled. This example focuses\u00a0on the p subshell, which fills from boron to neon.<\/p>\n

    \n
  • \n
    B (Z=5) configuration: 1s2<\/sup> 2s2<\/sup> 2p1<\/sup>\n\t\t\t\t<\/div>\n<\/li>\n
  • \n
    C (Z=6)\u00a0configuration:1s2<\/sup> 2s2<\/sup> 2p2<\/sup>\n\t\t\t\t<\/div>\n<\/li>\n
  • \n
    N (Z=7)\u00a0configuration:1s2<\/sup> 2s2<\/sup> 2p3<\/sup>\n\t\t\t\t<\/div>\n<\/li>\n
  • \n
    O (Z=8)\u00a0configuration:1s2<\/sup> 2s2<\/sup> 2p4<\/sup>\n\t\t\t\t<\/div>\n<\/li>\n
  • \n
    F (Z=9)\u00a0configuration:1s2<\/sup> 2s2<\/sup> 2p5<\/sup>\n\t\t\t\t<\/div>\n<\/li>\n
  • \n
    Ne (Z=10)\u00a0configuration:1s2<\/sup> 2s2<\/sup> 2p6<\/sup>\n\t\t\t\t<\/div>\n<\/li>\n<\/ul>\n

    Exceptions
    \n\t\t\t\t<\/em><\/h4>\n

    Although the\u00a0Aufbau rule\u00a0accurately predicts\u00a0the electron configuration of most elements, there are notable exceptions\u00a0among the transition metals and heavier elements. The reason these exceptions occur is that some elements are more stable with fewer\u00a0electrons in some subshells and more electrons in others (Table 1).
    \nTable 1<\/strong>: Exceptions to Electron Configuration Trends<\/em>
    \n\t\tPeriod 4:<\/strong>\u00a0\u00a0\u00a0\u00a0Period 5:<\/strong>
    \n\t\tChromium: Z:24 [Ar] 3d5<\/sup>4s1<\/sup>\u00a0\u00a0\u00a0\u00a0Niobium: Z:41 [Kr] 5s1<\/sup> 4d4<\/sup>
    \n\t\tCopper: Z:29 [Ar] 3d10<\/sup>4s1<\/sup>\u00a0\u00a0\u00a0\u00a0Molybdenum: Z:42 [Kr] 5s1<\/sup> 4d5<\/sup>
    \n\t\t\u00a0\u00a0\u00a0\u00a0\u00a0Ruthenium: Z:44 [Kr] 5s1<\/sup> 4d7<\/sup>
    \n\t\t\u00a0\u00a0\u00a0\u00a0\u00a0Rhodium:\u00a0 Z:45 [Kr] 5s1<\/sup> 4d8<\/sup>
    \n\t\t\u00a0\u00a0\u00a0\u00a0\u00a0Palladium: Z:46 [Kr] 4d10<\/sup>
    \n\t\t\u00a0\u00a0\u00a0\u00a0\u00a0Silver: Z:47 [Kr] 5s1<\/sup> 4d10<\/sup>
    \n\t\tPeriod 6:<\/strong>\u00a0\u00a0\u00a0\u00a0Period 7:<\/strong>
    \n\t\tLanthanum: Z:57 [Xe] 6s2<\/sup> 5d1<\/sup>\u00a0\u00a0\u00a0\u00a0Actinium: Z:89 [Rn] 7s2<\/sup> 6d1<\/sup>
    \n\t\tCerium: Z:58 [Xe] 6s2<\/sup> 4f1<\/sup> 5d1<\/sup>\u00a0\u00a0\u00a0\u00a0Thorium: Z:90 [Rn] 7s2<\/sup> 6d2<\/sup>
    \n\t\tGadolinium: Z:64 [Xe] 6s2<\/sup> 4f7<\/sup> 5d1<\/sup>\u00a0\u00a0\u00a0\u00a0Protactium: Z:91 [Rn] 7s2<\/sup> 5f2<\/sup> 6d1<\/sup>
    \n\t\tPlatinum: Z:78 [Xe] 6s1<\/sup> 4f14<\/sup> 5d9<\/sup>\u00a0\u00a0\u00a0\u00a0Uranium: Z:92 [Rn] 7s2<\/sup> 5f3<\/sup> 6d1<\/sup>
    \n\t\tGold: Z:79 [Xe] 6s1<\/sup> 4f14<\/sup> 5d10<\/sup>\u00a0\u00a0\u00a0\u00a0Neptunium: Z:93 [Rn] 7s2<\/sup> 5f4<\/sup> 6d1<\/sup>
    \n\t\t\u00a0\u00a0\u00a0\u00a0\u00a0Curium: Z:96 [Rn] 7s2<\/sup> 5f7<\/sup> 6d1<\/sup>
    \n\t\t\u00a0\u00a0\u00a0\u00a0\u00a0Lawrencium: Z:103 [Rn] 7s2<\/sup> 5f14<\/sup> 7p1<\/sup><\/p>\n

    Writing Electron Configurations
    \n<\/h3>\n

    When writing an electron configuration,\u00a0first write the energy level<\/strong> (the period), then the subshell<\/strong> to be filled and the superscript<\/strong>, which is the number of electrons in that subshell. The total number of electrons is the atomic number, Z. The rules above allow one to write\u00a0the electron configurations for all the elements in the periodic table.
    \nThree methods are used to write electron configurations:<\/p>\n

      \n
    1. \n
      orbital diagrams\n<\/div>\n<\/li>\n
    2. \n
      spdf notation\n<\/div>\n<\/li>\n
    3. \n
      noble gas notation\n<\/div>\n<\/li>\n<\/ol>\n

      Each method has its own purpose and each has its own drawbacks.<\/p>\n

      Orbital Diagrams
      \n\t\t\t\t<\/em><\/h4>\n

      An\u00a0orbital diagram, like those shown above,\u00a0is a visual way to reconstruct the electron configuration by showing each of the separate orbitals and the spins on the electrons. This is done by first determining the subshell (s,p,d, or f) then drawing in each electron according to the stated rules above.
      \nExample 4: Aluminum and Iridium
      \nWrite the electron configuration for aluminum and iridium.
      \nSOLUTION<\/strong>
      \n\t\tAluminum is in the 3rd period and it has an atomic number of Z=13. If we look at the periodic table we can see that its in the p-block as it is in group 13. Now we shall look at the orbitals it will fill: 1s, 2s, 2p, 3s, 3p. We know that aluminum completely fills the 1s, 2s, 2p, and 3s orbitals because mathematically this would be 2+2+6+2=12.\u00a0 The last electron is in the 3p orbital. Also another way of thinking about it is that as you move from each orbital block, the subshells become filled as you complete each section of the orbital in the period. The block that the atom is in (in the case for aluminum: 3p) is where we will count to get the number of electrons in the last subshell (for aluminum this would be one electron because its the first element in the period 3 p-block). This gives the following:
      \n\"\"\/
      \n\t\tNote that in the orbital diagram, the two opposing spins of the electron can be visualized. This is why it is sometimes useful to think about electron configuration in terms of the diagram. However,\u00a0because it is the most time consuming method, it is more common to write or see electron configurations in spdf notation and noble gas notation. Another example is the electron configuration of iridium:
      \n\"\"\/\"\"\/\"\"\/
      \n\t\tThe electron configuration of iridium is much longer than aluminum. Although\u00a0drawing out each orbital may prove to be helpful in determining unpaired electrons, it is very time consuming and often not as practical as the spdf notation, especially for atoms with much longer configurations. Hund’s\u00a0rule is also followed, as each electron fills up each 5d orbital before being forced to pair with another electron.<\/p>\n

      spdf Notation
      \n\t\t\t\t\t<\/em><\/strong><\/h4>\n

      The most common way to describe electron configurations is to write distributions in the spdf notation. Although\u00a0the distributions of electrons in each orbital are not as apparent\u00a0as in the diagram, the total number of electrons in each energy level is described by a superscript that follows the relating energy level. To write the electron configuration of an atom, identify the\u00a0energy level of interest\u00a0and write the number of electrons in the energy level as its superscript as follows: 1s2<\/sup>. This is the electron configuration of helium; it denotes a full s orbital.\u00a0The periodic table is used as a reference to accurately write the electron configurations of all atoms.\u00a0
      \nExample 5: Yttrium
      \nWrite the electronic configuration of Yttrium.
      \nSOLUTION<\/strong>
      \n\t\tStart with the straightforward\u00a0problem of\u00a0finding the electron configuration of the element yttrium. As always,\u00a0refer to the periodic table. The element yttrium (symbolized Y) is a transition metal, found in the fifth period and in Group 3. In total it has thirty-nine electrons. Its electron configuration is as follows:
      \n\u00a01s2<\/sup> 2s2<\/sup> 2p6<\/sup> 3s2<\/sup> 3p6<\/sup> 4s2<\/sup> 3d10<\/sup> 4p6<\/sup> 5s2<\/sup> 4d1<\/sup><\/strong>
      \n\t\tThis is a much simpler and more efficient way to portray electron configuration of an atom. A logical way of thinking about it is that all that is required is to fill orbitals\u00a0across a period and through orbital blocks. The number\u00a0of elements in each block is the same as in the energy level it corresponds. For example, there are 2 elements in the s-block, and 10 elements in the d-block. Moving across,\u00a0simply count how many elements fall in each block. Yttrium\u00a0is the first element in the fourth period d-block;\u00a0thus there is\u00a0one electron in that energy level. To check the answer, verify that the subscripts add up to the atomic number. In this case, 2+2+6+2+6+2+10+6+2+1= 39 and Z=39, so the answer is correct.
      \nA slightly more complicated example is the electron configuration of bismuth (symbolized Bi,\u00a0with Z = 83). The periodic table gives\u00a0the following electron configuration:
      \n1s2<\/sup> 2s2<\/sup> 2p6<\/sup> 3s2<\/sup> 3p6<\/sup> 4s2<\/sup> 3d10<\/sup>\u00a04p6<\/sup>5s2<\/sup> 4d10<\/sup> 5p6<\/sup> 6s2<\/sup> 4f14<\/sup> 5d10<\/sup> 6p3<\/sup><\/strong>
      \n\t\tThe reason why this electron configuration seems more complex is that the f-block, the Lanthanide series, is involved. Most students who first learn electron configurations often have trouble with configurations that must pass through the f-block because they often overlook this break in the table and skip that energy level. Its important to remember that when passing the 5d and 6d energy levels that one must pass through the f-block lanthanoidand actinoidseries. Keeping\u00a0this in mind, this “complex” problem is greatly simplified.
      \nAnother method (but less commonly used) of writing the spdf notation is the expanded notation format. This is the same concept as before, except that each individual orbital is represented with a subscript. The p, d, and f orbitals have different sublevels. The p orbitals are px,py<\/em>, and pz, <\/em>and if represented on the 2p energy with full orbitals would look like: 2px<\/sub>2<\/sup>\u00a02py<\/sub>2<\/sup> 2pz<\/sub>2<\/sup>. The expanded notation for neon (Ne, Z=10) is written as follows:
      \n1s2<\/sup> 2s2<\/sup>\u00a02px<\/sub>2<\/sup>\u00a02py<\/sub>2<\/sup>\u00a02pz<\/sub>2<\/sup><\/strong>
      \n\t\tThe individual orbitals are represented,\u00a0but the spins on the electrons are not; opposite spins are assumed. When representing the configuration of an atom with half filled orbitals, indicate\u00a0the two half filled orbitals. The expanded notation for carbon is written as follows:
      \n1s2<\/sup>\u00a02s2<\/sup>\u00a02px<\/sub>1<\/sup>\u00a02py1<\/sub><\/strong>
      \n\t\tBecause\u00a0this form of the spdf notation is not typically\u00a0used, it is not as important to dwell on this detail as it is to understand how to use the general spdf notation. \u00a0<\/p>\n

      Noble Gas Notation
      \n<\/h4>\n

      This brings up an interesting point about elements and electron configurations. As the p subshell is filled in the above example about the Aufbau\u00a0principle\u00a0(the trend from boron to neon), it reaches the group commonly known as the noble gases. The noble gases have the most stable electron configurations,\u00a0and are known for being relatively inert. All noble gases have their subshells filled and can be used\u00a0them as a shorthand way of writing electron configurations for subsequent atoms. This method of writing configurations is called the noble gas notation, in which the noble gas in the period above the element that is being analyzed is used to denote the subshells that element has filled and after which the valence electrons (electrons filling orbitals in the outer most shells) are written. This looks slightly different from spdf notation, as the reference noble gas must be indicated.
      \nExample 6: Vanadium
      \nWhat is the electronic configuration of vanadium (V, Z=23)?
      \nSOLUTION<\/strong>
      \n\t\tVanadium is the transition metal\u00a0in the fourth period and the fifth group. The noble gas preceding it is argon\u00a0(Ar, Z=18), and knowing that vanadium has filled those orbitals before it,\u00a0argon is used as the\u00a0reference noble gas. The noble gas in the configuration is denoted E, in brackets: [E]. To find the valance electrons that follow, subtract\u00a0the atomic numbers: 23 – 18 = 5. Instead of 23 electrons to distribute in orbitals, there are\u00a05. Now there is\u00a0enough information to write the electron configuration:
      \nVanadium, V: [Ar] 4s2<\/sup> 3d3<\/sup><\/strong>
      \n\t\tThis method streamlines the process of distributing electrons by showing the valence electrons, which determine\u00a0the chemical properties of atoms. In addition, when determining the number of unpaired electrons in an atom, this method allows quick\u00a0visualization of\u00a0the configurations of the valance electrons. In the example above, there are\u00a0a full s orbital and three half filled d orbitals.<\/p>\n

      Ions: Atoms with an Electrical Charge
      \n<\/h2>\n

      Atoms (or groups of atoms) in which there are unequal numbers of protons and electrons are called ions<\/em>. Usually, the number of protons and electrons in atoms are equal. But there are cases in which an atom can acquire an electrical charge.<\/p>\n

      An ion example
      \n<\/h2>\n

      For example, in the compound sodium chloride \u2014 table salt \u2014 the sodium atom has a positive charge and the chlorine atom has a negative charge.
      \nThe neutral sodium atom has 11 protons and 11 electrons, which means it has 11 positive charges and 11 negative charges. Overall, the sodium atom is neutral, and it’s represented like this: Na. But the sodium ion contains one more positive charge than negative charge, so it’s represented like this:
      \n\"\"\/
      \n\t\tThis unequal number of negative and positive charges can occur in one of two ways: An atom can gain a proton (a positive charge) or lose an electron (a negative charge).<\/p>\n

      Cations and anions
      \n<\/h2>\n

      So which process is more likely to occur? In general, it’s easy to gain or lose electrons but very difficult to gain or lose protons. So atoms become ions by gaining or losing electrons. And ions that have a positive charge are called cations<\/em>.
      \nThe progression goes like this: The sodium ion shown above is formed from the loss of one electron. Because it lost an electron, it has more protons than electrons, or more positive charges than negative charges, which means it’s now called the:
      \n\"\"\/
      \n\t\tLikewise, when the neutral magnesium atom loses two electrons, it forms the:
      \n\"\"\/
      \n\t\tNow consider the chlorine atom in sodium chloride. The neutral chlorine atom has acquired a negative charge by gaining an electron. Because it has unequal numbers of protons and electrons, it’s now an ion. And because ions that have a negative charge are called anions<\/em>, it’s now called the:
      \n\"\"\/<\/p>\n

      Other details about ions
      \n<\/h2>\n

      Here are some extra tidbits about ions:<\/p>\n

        \n
      • \n
        You can write electron configurations and energy level diagrams for ions. The neutral sodium atom (11 protons) has an electron configuration of:\n<\/div>\n

        \"\"\/
        \n\t\t\t\tThe sodium cation has lost an electron \u2014 the valence electron, which is farthest away from the nucleus (the 3s electron, in this case). The electron configuration of the sodium ion is:
        \n\"\"\/\n\t\t\t\t<\/li>\n

      • \n
        The electron configuration of the chloride ion is:\n<\/div>\n

        \"\"\/
        \n\t\t\t\tThis is the same electron configuration as the neutral Argon atom. If two chemical species have the same electron configuration, they’re said to be isoelectronic<\/em>.\n<\/li>\n

      • \n
        The preceding examples are all monoatomic (one atom) ions. But polyatomic (many atom) ions do exist. The ammonium ion is a polyatomic ion, or, specifically, a polyatomic cation. It is written as:\n<\/div>\n

        \"\"\/
        \n\t\t\t\tThe nitrate ion, is also a polyatomic ion, or, specifically, a polyatomic anion. It is written as
        \n\"\"\/\n\t\t\t\t<\/li>\n

      • \n
        Ions are commonly found in a class of compounds called salts, or ionic solids. Salts, when melted or dissolved in water, yield solutions that conduct electricity.\n<\/div>\n

        A substance that conducts electricity when melted or dissolved in water is called an electrolyte<\/em>. Table salt \u2014 sodium chloride \u2014 is a good example.
        \nOn the other hand, when table sugar (sucrose) is dissolved in water, it becomes a solution that doesn’t conduct electricity. So sucrose is a nonelectrolyte.
        \nWhether a substance is an electrolyte or a nonelectrolyte gives clues to the type of bonding in the compound. If the substance is an electrolyte, the compound is probably ionically bonded. If it’s a nonelectrolyte, it’s probably covalently bonded.\n<\/li>\n<\/ul>\n

        \u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0EVALUATION
        \n<\/strong>1.Write the electronic configuration of the first twenty elements.
        \n<\/strong>2.Write out the characteristics of the first three fundamental particles in an atom.
        \n<\/strong>3.An atom of an element is represented by X. How many electrons,protons and neutrons are in the atom? Write the electronic structure of the atom.
        \n<\/strong>
        \n\u00a0
        \n\u00a0
        \n\t\t\t<\/strong>\u00a0<\/p>\n","protected":false},"excerpt":{"rendered":"

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